GENERAL CHEMISTRY
122
LECTURE NOTES
Introduction
• Syllabus
Be familiar with and keep!!
• Calculators
Must purchase and use TI-30Xa or
updated nonprogrammable version !!
• Getting Help
Office Hours or by appointment
Help Sessions
• My Teaching Philosophy
1. It is absolutely necessary to have PP
slides in advance. You will also need to
takes notes!!!! Problem solving not
shown on notes.
2. Help you learn and really understand
some fundamentals of chemistry.
3. Help you learn to apply these
fundamentals to do problem solving.
4. Coach
5. Athlete Analogy
• Study Tips
1. Amount of time required depends
on background and abilities.
2. Practice problems / homework until
you can do without text or example.
3. Review / study prior to class.
4. Must understand concepts and be
able to apply to problem solving.
5. Get help when needed!!
ORGANIC COMPOUNDS
• Recognition of Structures: (see handout)
Alkanes
Ketones
Alkenes
Carboxylic Acids
Alkynes
Amines
Alcohols
Benzene
Ethers
Phenyl group
Aldehydes
CHAPTER 11
LIQUIDS, SOLIDS, AND
INTERMOLECULAR
FORCES
PROPERTIES OF MATTER
I. KINETIC MOLECULAR THEORY
* All matter is composed of tiny particles
(atoms or molecules) in constant motion.
* Increasing temperature increases the
motion of the particles.
Can explain properties of matter that we observe
on the macroscale on the basis of the behavior of
the molecules on the nanoscale.
II. INTERMOLECULAR FORCES
Forces primarily responsible for differences
between (solids, liquids) and gases.
What are they?
1. Attractive forces between molecules
2. They are not covalent bonds.
3. Much weaker than covalent bonds.
Significant in Understanding Properties of
Matter!!!
Covalent Bonds vs. Intermolecular Forces
What are they?
1.
London Forces
Noncovalent interaction between all
molecules
Due to induced dipole moments created
by movement of electron clouds.
London Dispersion Forces are the only
Intermolecular Force between nonpolar
molecules.
Factors Affecting Their Strength
Greater Polarizability Stronger London
Dispersion Forces
*** Larger atoms or molecules
Stronger Dispersion (London forces)
(approximated to atomic or molecular mass)
*** Molecular Shape  Stronger London Forces
for elongated vs. compact molecules.
Which has strongest London Forces?
a. Ne He Ar
b. F2 Cl2 Br2
Significance?
I2
Which has the highest boiling point?
(Stronger Intermolecular Forces 
Higher Melting and Boiling Points)
a. n-Pentane (C5H12) vs. n-Octane (C8H18)
b. n-Octane vs. Isooctane
(see next slide)
2. Dipole-Dipole Forces
Noncovalent interaction between polar
molecules or groups.
Attractive force between a partially positive
region of one molecule in close proximity to a
partially negative region of another molecule
The greater the polarity of a molecule, the
stronger the dipole-dipole forces.
Which has strongest Dipole-dipole forces?
Why?
H-F vs H-Br ??
Examples
1. Which types of molecules commonly have
stronger intermolecular forces? Why?
polar vs. nonpolar molecules
2. Which has strongest dipole-dipole forces?
propane (C3H8) or acetaldehyde (CH3CHO)
3. Describe / show dipole-dipole interactions for
Br-Cl molecules.
3. Hydrogen Bonding
Special dipole-dipole interaction.
Partially positive H atom which is covalently
bonded to an electronegative atom
(O,N,F) has an attractive force for
another electronegative atom (O,N,F).
Greater # H bonds  greater intermolecular
attractive forces
Hydrogen Bonding is strongest of the three
Intermolecular Forces !!
Examples:
1. Can they hydrogen bond?
1. Water?
2. Methanol (CH3OH)
3. Methane (CH4)
4. Ammonia (NH3)
2. Which forms greatest degree of
hydrogen bonding? ethanol (CH3CH2OH)
or acetic Acid (CH3CO2H)
Properties of Matter Related
To Intermolecular Forces
1. Solubility Rule: “Likes Dissolve Likes”
Polar and ionic compounds are soluble in
polar solvents.
Nonpolar compounds are soluble in
nonpolar solvents.
Is solid KBr more soluble in water or gasoline?
Is vegetable oil more soluble in water or gasoline?
2. Greater Intermolecular Forces Related
to Higher Melting and Boiling Points
3. Unusual Properties of Water
**Primarily due to ability to form several “H” bonds.
Water thus forms strong intermolecular forces.
1. Low mass, yet liquid at room temperature.
2. High specific heat capacity
3. High heat of vaporization, high boiling
point.
4. Ice less dense than liquid water. Ice floats.
Why?
PROPERTIES OF LIQUIDS RELATED
TO INTERMOLECULAR FORCES
I. VISCOSITY
The resistance of a liquid to flow.
The stronger the intermolecular
forces, the greater the viscosity.
Viscosity decreases as temp. increases.
II. SURFACE TENSION
* Intermolecular attractions between molecules
of a liquid create surface tension.
(unlike a gas)
* Uneveness of the forces on the surface causes
the surface of the liquid to contract.
* Surface tension = energy required to increase
the surface area by a unit amount.
* Stronger intermolecular forces stronger or
higher surface tension.
III. CAPILLARY ACTION
*
Molecules of liquid can interact with the
molecules of the container.
*
Must consider attractive forces between liquid
molecules compared to attractive forces
between liquid and container.
*
What is capillary action?
*
What is a meniscus??
IV. VAPOR PRESSURE
*
Volatility – tendency of a liquid to vaporize
*
Vaporization- when a molecule moves from
the liquid phase to the gas (vapor) phase.
Why / how does this happen? (see next Figure)
a. Liquid molecules have varying kinetic energy
b. If kinetic energy of molecules in liquid overcomes
intermolecular forces in liquid, they can escape to
gas.
c. Increased temp. increased vaporization Why?
* (Equilibrium) Vapor Pressure
In a closed container the liquid and gas phases of a
substance come to dynamic equilibrium.
What does that mean?
The pressure of the gas (vapor) above a liquid at
equilibium is called the equilibrium vapor pressure.
What is creating the pressure?
Higher volatilitymore molecules in gas phase
higher vapor pressure
Example Vapor Pressure Problem
Equilibrium is established between a small
quantity of CCl4(l) and its vapor at 400C in a flask
having a volume of 285 mL. The total mass of
vapor present is 0.480 g. What is the vapor
pressure of CCl4, in mm Hg, at 400C?
IV. Boiling Point of Liquid
* Boiling Point- temperature at which the
(equilibrium) vapor pressure equals the
atmospheric pressure.
* Normal BP – when atmospheric pressure = 1 atm.
Higher intermolecular forces lower
vapor pressure higher boiling point
How does size of molecules affect BP?
1. Which has highest BP?
Methane (CH4)
Ethane
(C2H4)
Propane
(C3H6)
2. Which has highest BP?
Br2
H2O
How does change in atmospheric pressure
affect BP?
PHASE CHANGES OF MATTER
Phase Change
Name
SolidLiquid
Liquid Solid
Liquid Gas
Gas Liquid
Solid Gas
Gas Solid
Melting, Fusion
Freezing, Crystallization
Vaporization
Condensation
Sublimation
Deposition
I. PHASE TRANSITIONS
a. Raindrops hit cold metal surface on car
and it becomes covered with ice.
b. Frozen clothes on line dry at below
freezing temperatures (of H2O.)
c. Rubbing alcohol spilled on the palm of
the hand feels cool.
What must be supplied or removed for phase changes?
II. Enthalpy of Phase Transitions
Enthalpy (ΔH) = heat energy change under
constant pressure and “temperature”
conditions.
If “-”  heat produced
If “+”  heat required
Liquid ⇌ Gas
Δ H vaporization = - ΔHcondensation
H2O (l)  H2O (g)
ΔHvap = + 40.7 kJ/mol
H2O (g)  H2O (l)
ΔHcond = - 40.7 kJ/mol
Solid  Liquid  Gas
Endothermic processes
Gas  Liquid  Solid
Exothermic processes
** Consider problem solving using ΔH !!
III. Problems Using Enthalpy:
Isopropyl alcohol, C3H7OH (60.0 g/mol), is
used in rubbing alcohol mixtures. Alcohol
on the skin cools by evaporation. How
much heat is absorbed by the alcohol if 10.0
g evaporate? The enthalpy of vaporization
for isopropyl alcohol is 42.1 kJ/mol.
Note: Problems with changes in
temperature of a substance which
also includes a phase transition.
How much energy is required to heat 15g of
water from –10oC to 60oC?
Δ H fusion = 6.020 kJ/mol
Δ H vaporization = 40.7 kJ/mol
Specific Heat Capacity solid H2O = 2.06 J/goC
Specific Heat Capacity liquid H2O = 4.184 J/goC
Specific Heat Capacity gas H2O = 2.10 J/goC
See next Figure or Figure 11.36 in text
IV. Phase Diagrams
* Graphical representation to summarize
conditions (pressure and temperature)under
which different states of a substance are
stable. See Fig 11.14
* Be familiar with:
a. What the lines represent
b. Identify what states are present at some T and P
c. Terms used
Terms Used:
1.
Triple Point - P and T where all 3 phases exist in
equilibrium.
2.
Critical Point – endpoint of line separating liquid
and gas. At this point liquid and gas are
indistinguishable. Occurs at critical temperature
and critical pressure. Explain
3.
Supercritical Fluid - substance that exists above
critical temperature and pressure. Has properties
of both liquid and gas. Significance
PHASE DIAGRAM FOR WATER
PHASE DIAGRAM FOR CO2
SOLID STATE INFORMATION
I. Classified as Amorphous or Crystalline
Amorphous Solids - a solid that has a disordered
structure. No well defined arrangement of basic units
(atoms, molecules, or ions) at nanoscale level.
Examples – cement, glass, optical fibers.
Crystalline Solids - a solid that has an ordered
structure. Well defined symmetrical arrangement of
basic units (atoms, molecules, or ions). Composed of
one or more crystals with well defined 3-D structure.
II. Classification of Crystalline Solids
A. Ionic Solid –
a solid that consists of positive ions (cations) and
negative ions (anions) held together in a lattice by the
electrical attraction of the opposite charges (ionic
bonds). Very strong bonds. NaCl
B. Molecular Solid –
a solid that consists of atoms or molecules held
together by intermolecular forces (London forces,
dipole-dipole forces, and hydrogen bonds). Usually
involves nonmetals. I2, H2O
C. Atomic Solids
1. Metallic Solid – a solid that consists of
positive cores of metal atoms held together
by a surrounding sea of electrons (metallic
bond). Good electrical conductors. Fe
2. Network Solid - a solid that consists of
atoms held together in large networks or
chains by covalent bonds. Network of
nonmetal atoms. Graphite, Diamond
Crystal Structure of Graphite
Hexagons of sp2hybridized
carbon atoms.
Forces between
layers are
relatively weak.
Crystal Structure of Diamond
Three-dimensional
network is extremely
strong, rigid.
What kind of forces
must be overcome to
melt diamond?