Chemistry II Syllabus
I. Structure of Matter
A.
Atomic theory and atomic structure
1.
Evidence for the atomic theory
2.
Atomic masses; determination by chemical and physical means
3.
Atomic number and mass number; isotopes
4.
Electron energy levels: atomic spectra, quantum numbers, atomic
orbitals
5. Periodic relationships including, for
example, atomic radii, ionization energies, electron affinities, oxidation
states
B.
Chemical bonding
1.
Binding forces
a. Types: ionic, covalent, metallic,
hydrogen bonding, van der Waals (including London
dispersion forces)
b. Relationships to states, structure, and
properties of matter
c. Polarity of bonds, electronegativities
2.
Molecular models
a. Lewis structures
b. Valence bond: hybridization of orbitals,
resonance, sigma and pi bonds
c. VSEPR
3. Geometry of molecules and ions,
structural isomerism of simple organic molecules and coordination complexes;
dipole moments of molecules; relation of properties to structure
C.
Nuclear chemistry: nuclear equations, half-lives, and radioactivity; chemical
applications
II. States of Matter
A.
Gases
1.
Laws of ideal gases
a. Equation of state for an ideal gas
b. Partial pressures
2.
Kinetic molecular theory
a. Interpretation of ideal gas laws on the
basis of this theory
b. Avogadros hypothesis and the mole
concept
c. Dependence of kinetic energy of molecules
on temperature
d. Deviations from ideal gas laws
B.
Liquids and solids
1.
Liquids and solids from the kinetic-molecular viewpoint
2.
Phase diagrams of one-component systems
3.
Changes of state, including critical points and triple points
4.
Structure of solids; lattice energies
C.
Solutions
1.
Types of solutions and factors affecting solubility
2.
Methods of expressing concentration (use of normalities
is not tested)
3.
Raoults law and colligative
properties (nonvolatile solutes); osmosis
4.
Nonideal behavior (qualitative aspects)
III. Reactions
A.
Reaction types
1. Acid-base reactions; concepts of
Arrhenius, Brønsted-Lowry, and Lewis; coordination
complexes; amphoterism
2.
Precipitation reactions
3.
Oxidation-reduction reactions
a. Oxidation number
b. The role of the electron in
oxidation-reduction
c. Electrochemistry: electrolytic and
galvanic cells; Faradays laws; standard half-cell potentials; Nernst equation;
prediction of the direction of redox reactions
B.
Stoichiometry
1.
Ionic and molecular species present in chemical systems: net ionic equations
2.
Balancing of equations, including those for redox
reactions
3. Mass and volume relations with emphasis
on the mole concept, including empirical formulas and limiting reactants
C.
Equilibrium
1. Concept of dynamic equilibrium, physical
and chemical; Le Chateliers principle;
equilibrium constants
2.
Quantitative treatment
a. Equilibrium constants for gaseous reactions:
Kp, Kc
b. Equilibrium constants for reactions in
solution
(1) Constants for acids and bases; pK; pH
(2) Solubility product constants and their
application to precipitation and the dissolution of slightly soluble compounds
(3) Common ion effect; buffers; hydrolysis
D.
Kinetics
1.
Concept of rate of reaction
2. Use of experimental data and graphical
analysis to determine reactant order, rate constants, and reaction rate laws
3.
Effect of temperature change on rates
4.
Energy of activation; the role of catalysts
5.
The relationship between the rate-determining step and a mechanism
E.
Thermodynamics
1.
State functions
2. First law: change in enthalpy; heat of
formation; heat of reaction; Hesss law; heats of vaporization and fusion; calorimetry
3. Second law: entropy; free energy of
formation; free energy of reaction; dependence of change in free energy on
enthalpy and entropy changes
4.
Relationship of change in free energy to equilibrium constants and electrode
potentials
IV. Descriptive Chemistry
Knowledge
of specific facts of chemistry is essential for an understanding of principles
and concepts. These descriptive facts, including the chemistry involved in
environmental and societal issues will be taught throughout the course to
illustrate and illuminate the principles of chemistry. The following areas
should be covered:
1.
Chemical reactivity and products of chemical reactions
2. Relationships in the periodic table:
horizontal, vertical, and diagonal with examples from alkali metals, alkaline
earth metals, halogens, and the first series of transition elements
3. Introduction to organic chemistry:
hydrocarbons and functional groups (structure, nomenclature, chemical
properties)
V. Laboratory
Students
will acquire these skills in the laboratory:
making observations of chemical reactions and
substances
recording data
calculating and interpreting results based on the
quantitative data obtained
communicating effectively the results of experimental work
VI. Chemical Calculations
The
following list summarizes types of problems either explicitly or implicitly
included in the preceding material. Attention will be given to significant
figures, precision of measured values, and the use of logarithmic and
exponential relationships. Critical analysis of the reasonableness of results
will be encouraged.
1.
Percentage composition
2.
Empirical and molecular formulas from experimental data
3.
Molar masses from gas density, freezing-point, and
boiling-point measurements
4.
Gas laws, including the ideal gas law, Daltons law, and Grahams law
5.
Stoichiometric relations using the concept of the mole; titration calculations
6.
Mole fractions; molar and molal solutions
7.
Faradays laws of electrolysis
8.
Equilibrium constants and their applications, including their use for
simultaneous equilibria
9.
Standard electrode potentials and their use; Nernst equation
10.
Thermodynamic and thermochemical calculations
Classwork, Tests, and Grades
ABCI--You
CAN fail if YOU allow it to happen, but you will never receive a D or an F
grade on any work you do. Work
that is below the 75% standard will be marked I for incomplete, and will need
to be brought up to standards before any grade is assigned to your work for the
course. Doing substandard work in
class, just like in the work world, is not acceptable.
Practice
Work--Even though they may not included as part of your grade,
practice assignments, homework, reading activities, etc., must be completed at
or above the 75% level. If they
are missing, or poorly done, they will receive a mark of I, and must be
redone or completed. Is in the
grade book need to be made up before any grade is assigned for the course.
Incompletes become Fs two weeks after the end of the quarter.
Tests—Most
written tests are given in two parts.
The first part assesses your achievement measured against the minimum
standards for the unit. If that is
the only part of the test that you do, the best grade that can be earned is a
C.
If you want to show that your achievement is worthy of a B, you
must meet the 75% minimum on the first part of the test, and score at or above
50% on the advanced part of the test.
If you want to demonstrate A achievement,
you must meet the 75% minimum on the first part of the test, and score at or
above 75% on the advanced part of the test.
All scores can be improved by demonstrating that you have met the
goals, either by taking a retest, repairing the earlier test, or completing
an alternative assessment. These
can be done only after additional teaching/learning has taken place to remedy
any deficiencies. Since we all
learn differently, I will use my professional judgment as to which method will
be the most suitable in each instance.
Alternative
Assessments--On occasion, you will be asked to demonstrate your achievement
by writing a paper, completing a project, or doing some other formal
activity. Sometimes, you will be
given the opportunity to choose the assessment method that you will use. You will be given a rubric (checklist)
when the assignment is made so you will know exactly what is expected of you in
order to meet minimum standards.
Grades are earned on a percentage of total points possible.
A 95
A- 92
B+ 90
B 88
B- 85
C+ 82
C 78
C- 75
I <75
