Download 12550 College Chemistry 1998

Course Title: College Chemistry 1A / 1B
Small Assembly Session (SAS)
1A / 1B, and Laboratory 1A / 1B
Date Adopted:
Department: Science
UC / CSU Requirements:
High School-College / University--Transferable
Prerequisite: C or better in Chemistry
C or better in Algebra II
or concurrent enrollment
in Algebra II; Instructor’s approval
July 15, 1998
Fulfills CSF Requirement: Yes
Fulfills H / S Graduation Credit
as: Required ___ Elective _X__
Length Of Course: Two semesters
Semester Units (Credits):
High School--5 / 5 (10 per year)
College--Lecture--4 / 3
Small Assembly Session (SAS) 1A / 1B--0 / 1
Laboratory--1 / 1
Grade Level: 11 and 12
I. Course Description
College Chemistry is a comprehensive theoretical college level study of matter, the physical and
chemical changes that matter undergoes, and the relation between changes in composition and changes
in energy.
This course may be taken for high school and / or for college credit. If taken for college credit,
registration with Antelope Valley College is required along with payment of appropriate registration fees
(may be waived). College Chemistry is equivalent to Chemistry 1A / 1B, Small Assembly Session
(SAS) 1A /1B and Chemistry Laboratory 1A / 1B as taught at Antelope Valley College. Lecture and
SAS would be scheduled at the high school as part of the student’s regular daily program. Laboratory
sessions would be held one night per week (6 to 9 PM) at Antelope Valley College participating students
providing their own transportation to Antelope Valley College. College Chemistry is an alternative to
AP Chemistry. College credits earned would be transferable to other colleges or universities.
II. Rationale
This program will establish a cooperative chemistry program between Antelope Valley Union High
District and Antelope Valley College. As mentioned above, this cooperative effort involves using
Antelope Valley College’s laboratory facilities during the evening and instructing students during the
day at a high school location. Students desiring college credit would complete registration and pay
appropriate fee, if required, to Antelope Valley College.
Why this cooperative effort? Students enrolled in AP Chemistry, as is true with other advanced
placement courses, are required to complete a course in the academic area and to
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successfully complete a rigorous comprehensive examination only to be informed by some participating
colleges or universities that they will not accept the advanced placement course / scores for advanced
placement or that they will have to retake a portion of or to retake the entire course of study.
In Chemistry the problem is compounded by laboratory experimentation--many college and university
chemistry personnel believe that the laboratory sessions taught in high school AP Chemistry classes
don’t measure up to the same standards of scientific and academic rigor as do their college / university
counterparts.
By having students enrolled in an accredited college chemistry program, we believe credits earned would
be credits awarded and transferred.
III. Goals, Objectives, and Performance Indicators
I.0 GOAL: Students will gain an understanding of the Atomic Structure of Matter.
1.1 Obj: Students will develop an understanding of Atomic Theory of Matter and atomic
structure.
1.1.1: Students will determine atomic masses by mathematical and chemical and physical
means verifying the Atomic Theory of Matter.
1.1.2: Students will determine quantitatively the atomic number, mass number (atomic
mass), the number of basic subatomic particles (protons, neutrons, and electrons)
of element and isotopes of elements.
1.1.3: Students will construct according to modern Quantum Theory an atomic electron
distribution diagram, write elemental electron configurations, explain atomic
emission and adsorption spectra of elements qualitatively and quantitatively using
Planck’s Equation, assign quantum numbers (n, l, m and s) to each electron in
accordance with the Aufbau Principle, Pauli’s Exclusion Principle and Hund’s
Rule.
1.1.4: Students will identify the different periods and families, representative, transition
and inner transition elements of the Periodic Table and using the Periodic Law
able to predict the periodic properties (atomic radius, metallic character,
ionization energies, electronegativities, electron affinities, chemical behavior, etc.)
of the elements as a function of their atomic number.
1.1.5: Given the contributions of Aristotle, Democritus, Dalton, Mendeleev, Thomson,
Millikan, Rutherford, Chadwick, Moseley,Seaborg and others, students will relate
each contribution to the Periodic Table and view the Periodic Table as a
continuously evolving scientific tool.
1.1.6: Students will understand the formation of compounds from element and
differentiate between elements, mixture and compounds and heterogeneous and
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homogeneous mixtures.
1.1.7: Applying the mole concept and Avogadro’s number, students will stoichiometricly
calculate molar masses, moles, atoms, molecules of elements and compounds, and
empirical and molecular formulas, including verification of the chemical Laws of
Constant and Multiple Proportions.
1.2 Obj: Students will understand Chemical Bonding
1.2.1: Students will identify and differentiate between the binding forces between atoms
and the atoms of molecules (Ionic, covalent, metallic, hydrogen dipole and
dispersion forces), relate these bonding forces to state, structure, and the
properties of matter.
1.2.2: Using Valence Bond Theory, Valence Shell-Eelectron Pair Repulsion, Molecular
Orbital Theory, and the Octet Rule, including expanded octets, students will draw
Lewis electron dot and resonance structures, atomic and molecular orbital
diagrams, including hybridization of atomic orbitals, sigma and pi bonding.
1.2.3: Using Valence Bond Theory and Valence Shell-Eelectron Pair Repulsion VSEPR),
students will construct molecular models and identify the geometric shape of
molecules, geometric isomers and coordination compounds, predict the dipole
moment of molecules, relate these shapes to state, structure, and the properties of
matter.
1.2.4: Applying the Laws of Constant and Multiple Proportions, the Periodic Table, and a
knowledge of the rules for the systematic nomenclature of inorganic and organic
compounds, students will name and / or write the correct formula for Inorganic,
coordination, and elementary organic compounds, including identification of
common organic functional groups.
1.3 Obj: Students will develop knowledge of nuclear chemistry
1.3.1: Students will differentiate between the isotopes of an element, distinguish between
fission and fusion, balance nuclear transmutation equations, recognize that nuclear
transmutation reactions are first order reactions and calculate the half-life,
quantity remaining, and rate constant for these nuclear reactions.
1.3.2: Using the concept of mass defect, students will predict the stability of nuclei and
use Einstein’s equation for mass-energy to verify the Law of Conservation of
Mass-energy and apply these concepts to every day experiences (nuclear power,
medicine, weapons, etc.).
2.0 GOAL: Students will develop and knowledge of the States of Matter
2.1 Obj: Students will develop a knowledge of the Kinetic Theory of Matter and its application to
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the gaseous state.
2.1.1: Students will relate the Kinetic Theory of Matter to the gaseous state, use the Ideal
Gas Law, Combined Gas Law, Boyle's Law, Charles Law, Avogadro’s Law,
Dalton's Law, etc., to calculate and predict the qualitative and quantitative
behavior of gases to changes in pressure, temperature and volume.
2.1.2: Students will relate the Kinetic Theory of Matter to the Absolute Kelvin
Temperature Scale to predict the behavior of a gas to changing temperatures and
use Graham’s Law of Diffusion to calculate the relative velocities of molecules.
2.1.3: Students will predict the qualitative difference in the behavior of real gases
compared ideal gases and use van der Waal”s equation to calculate the
quantitative difference.
2.2 Obj: Students will develop a knowledge of the Kinetic Theory of Matter and its application to
the liquid and solid state.
2.2.1: Students will relate the Kinetic Theory of Matter to the liquid and solid state by
constructing heating and cooling diagrams and one component phase diagrams
and labeling these diagrams with the various states and critical and triple points;
and predict the state of matter present at various temperatures and pressures.
2.2.2: Students will explain qualitatively and calculate quantitatively the energy values
(vaporization, ionization, solidification) in the formation of a solid crystalline
lattice (lattice energy) using the Born-Haber cycle.
2.2.3: Students will quantitatively determine from calorimetric data and specific heats
Enthalphies of Reaction, Heats of Vaporization, and Heat of Fusion.
2.3 Obj: Students will develop an understanding of the importance of solution in chemistry.
2.3.1:Students will explain qualitatively and quantitatively using dissociation and
hydration equations the solution process and the factors affecting the solubility of
a substance.
2.3.2: Students will calculate the concentration of chemical solution using percentages by
mass, molarity, molality, normality,and formality, including dilution.
2.3.3: Students will explain qualitatively and calculate quantitatively the collegative
properties of boiling point elevation, freezing point depression, osmotic pressure,
Raoult’s Law as they relate to dilute solutions.
3.0 GOAL: Students will develop a knowledge of chemical reactions.
3.1 Obj: Students will develop a knowledge of the different types of chemical reactions.
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3.1.1: Students will write and balance by inspection molecular and net ionic equations
and classify chemical reactions as to type: synthesis, decomposition, combustion,
single displacement, double displacement acid-base, and double displacementprecipitation.
3.1.2: Students will predict the products of a variety of chemical reactions (synthesis,
decomposition, combustion, single displacement, double displacement acid-base,
and double displacement-precipitation) and write and balance chemical equation
illustrating their prediction.
3.1.3: Students will explain quantitatively and illustrate quantitatively using chemical
equations the different theories of acid-base chemistry: Arrhenius, BronstedLowry, and Lewis.
3.1.4: After assigning oxidation numbers, students will balance oxidation-reduction
equations using the half-reaction and the electron transfer methods and apply the
concept of oxidation-reduction to electrochemical and electrolytic cells, apply
Faraday’s Law of Electrolysis to electrolytic cells, use standard electrical voltages
and the Nernst equation to predict changes in standard electrical potentials and
direction of oxidation -reduction reactions.
3.2 Obj: Students will develop an appreciation for the application of stoichiometry to a variety of
chemical situations.
3.2.1: Using the mole method students will stoichiometricly calculate mole ratios, massmass, mass-volume, and volume-volume relationships from balanced chemical
equations involved in chemical and physical changes, including limiting reagents,
solution formation, precipitation, and titrations.
3.3 Obj: Students will develop an understanding of chemical equilibrium.
3.3.1: Students will explain qualitatively the dynamic nature of chemical equilibrium and
differentiate between equilibrium and complete chemical reactions.
3.3.2: Students will apply La Chatelier’s Principle to a system at equilibrium and predict
the response of the chemical system to an applied stress.
3.3.3: Students will apply the Law of Mass Action (Chemical Equilibrium) to a chemical
equation and write the equilibrium law expression for the chemical system and
qualitatively explain the meaning the equilibrium law constant.
3.3.4: Students will apply the law of chemical equilibrium a variety of chemical systems
and quantitatively calculate equilibrium constants and / or concentrations for
gaseous systems (Kp, Kc), acids and bases (Ka, Kb, Kw) including pH, pOH,
buffer, and solubilities (Ksp), including their application to precipitation,
solubility, and the common ion effect, and hydrolysis (Kh).
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3.4 Obj: Students will develop knowledge of reaction kinetics.
3.4.1: Students will qualitatively relate the Collision Theory of Reaction to the rate of a
chemical reaction and predict the effect of changing temperatures, pressures,
concentrations, and catalysts on the rate of a chemical reaction.
3.4.2: Students will use differential rate laws and experimental concentrations to
determine the order of reaction, the reaction rate law, and calculate the rate
constant.
3.4.3: Students will construct reaction coordinate diagrams and calculate from those
diagrams heats of reactions and activation energies and use the diagram to predict
the effect of a catalyst on the activation energy.
3.4.4: Students will qualitatively explain and quantitatively illustrate with chemical
equations the difference between a chemical reaction and a chemical mechanism
and determine from the mechanism of a chemical reaction the rate-determining
step and the reaction rate law.
3.5 Obj: Students will develop knowledge of chemical thermodynamics.
3.5.1: Students will qualitatively explain the First, Second, and Third Law of Chemical
Thermodynamics, why they are state functions, apply these laws to chemical
systems, and predict their response.
3.5.2: Using a Table of Standard Enthalphies of Formation, Standard Entropies of
Formation, and Standard Free Energies of Formation, and Hess’s Law students
will quantitatively calculate Enthalphies of Reaction, Entropies of Reaction and
Free Energies of Reaction.
3.5.3: Students will quantitatively relate the free energy, enthalphy, and entropy using two
to determine the third and further calculate effect of the free energy to the
equilibrium constant of a chemical reaction at equilibrium and the electrical
potential of a oxidation-reduction reaction.
4.0 Goal: Students will gain an understanding of the role mathematics plays in Chemistry.
4.1 Obj: Students will develop a knowledge of the application of mathematical skills to the study
of Chemistry
4.1.1: Students are expected to be able to demonstrate the abilities to perform the
following types of calculations, including the use of logarithmic and exponential
relationships. They are also expected to be aware of and pay attention to
significant figures, accuracy and precision of measurements.
a. Percentage calculations
b. Empirical and molecular formulas from experimental data
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c. Molar masses from gas densities, freezing point and boiling point
measurements
d. Gas laws, including the ideal gas law, Dalton’s law, and Graham’s law
e. Stoichiometry and limiting reagent stoichiometry using the mole method
f. Solution calculations including molarity and molarity
g. Faraday’s law of electrolysis
h. Equilibrium calculations and their applications
i. Standard electrode potentials and their uses; Nernst equation
j. Thermodynamic and thermochemical calculations; Hess’s Law
k. Kinetics calculations
IV. OUTLINE OF TOPICS:
I. Structure of Matter
A. Atomic theory and atomic structure.
1. Evidence for the atom theory
2. Determinations of atomic masses by chemical and physical means
3. Atomic number and mass number; isotopes
4. Electron energy levels, atomic spectra, quantum numbers, and atomic orbitals
5. Periodic properties of the elements.
6. Evolution of the periodic table of elements.
B. Chemical Bonding
1. Binding forces
a. Ionic, covalent, metallic, hydrogen dipole and dispersion forces
b. Relationship to states, structure and properties of matter
c. polarity of bonds and electronegativity
2. Molecular models
a. Lewis electron structures and resonance structures
b. Valence bond theory, atomic orbitals and hybridization of atomic orbitals,
sigma and pi bonds
c. VSEPR theory
3. Geometry of molecules and ions
a. Geometric shape of molecules and ions
b. Structural isomers of simple organic molecules and coordination compounds
c. Dipole momenta of molecules
d. Relation of properties to structure
4. Chemical nomenclature
a. Inorganic compounds
b. Coordination compounds
c. Elementary organic compounds and functional groups
C. Nuclear chemistry
1. Isotopes
2. Nuclear transmutation equations
3. Half-lives and radioactivity
4. Chemical application
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II. States of Matter
A. Gases
1. Laws of ideal gases
a. Equation of state for an ideal gas
b. Partial pressures
c. Graham’s law of diffusion
2. Kinetic-molecular theory
a. Interpretation of ideal gas laws on the basis of this theory
b. Avogadro’s hypothesis and the mole concept
c. Dependence of kinetic theory of molecules on temperature
d. Deviations from ideal gas laws
B. Liquids and solids
1. Liquids and solids form the kinetic-molecular theory viewpoint
2. Phase diagrams of one-component systems
3. Heating and cooling diagrams
4. Changes of state including critical points and triple points
5. Structure of solids and lattice energies
C. Solutions
1. Types of solutions and factors affecting solubility
2. Methods of expressing concentration
a. Percentage by mass
b. Molarity
c. Molality
d. Normality (optional)
e. Formality (optional)
3. Colligative properties
a. Freezing point depression
b. Boiling point elevation
c. Osmotic pressure
d. Raoult’s law
4. Non-ideal behavior (qualitative aspects)
III. Reactions
A. Reaction types
1. Acid-base reactions
a. Arrhenius
b. Bornsted-lowry
c. Lewis
d. Amphoterism
2. Coordination compounds
3. Precipitation reactions
4. Oxidation-reduction reactions
a. Oxidation numbers
b. Role of the electron in oxidation-reduction
c. Electrochemistry
-Electrolytic cells
-Galvanic cells
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-Faraday’s laws of electrolysis
-Standard half-cell potentials
-Nernst equations
- Prediction of the direction of redox reactions
B. Stoichiometry
1. Ionic and molecular species in chemical systems
2. Net ionic equations
3. Balancing equations--simple and oxidation-reduction
4. Mole concept
a. Molar mass and Avogadro’s number
b. Mass and volume stoichiometric calculations including limiting reagent
c. Empirical and molecular formula calculations
C. Equilibrium
1. Equilibrium
a. Dynamic nature of equilibrium
b. Physical and chemical equilibrium
c. Le Chaterler’s principle
d. Law of mass action and the equilibrium law expression
e. Equilibrium constants and their meaning
2. Quantitative aspects of equilibrium
a. Equilibrium constants for gaseous reactions, Kp, Kc
b. Equilibrium constants for reactions in solution
1. Acid-base equilibrium constants, Ka, Kb
2. Ion product constant for water, Kw
2. pH, pOH, buffers, and pK’s
c. Solubility product constants, Ksp--calculations and their application to
precipitation and the dissolution of slightly soluble compounds, common ion
effect on solubility
3. Hydrolysis constants, Kh
D. Kinetics
1. Concept of rate of reaction
2. Use of differential rate laws to determine the order of reaction and rate constant for
experimental data
3. Effect of temperature change on rates
4. Reaction coordinate diagram
5. Energy of Activation
6. Role of catalysts in reaction kinetics
7. Mechanisms of chemical reactions and rate determining steps
E. Thermodynamics
1. State functions
2. First law of thermodynamics
a. Changes in enthalpy
b. Heat of formations
c. Heats of reaction
d. Hess’s law
e. Heats of vaporization
f. Heats of fusion
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g. Calorimetry
3. Second law of thermodynamics
a. Entropy
b. Free energy of formation
c. Free energy of reaction
d. Relation between free energy, enthalpy and entropy
4. Relationship of change in free energy to equilibrium constants and electrode potentials
IV. Descriptive Chemistry
A. Chemical reactivity and products of chemical reactions
B. Classification of chemical reactions
1. Synthesis
2. Decomposition
3. Double displacement
a. Acid-base
b. Precipitation
5. Combustion
6. Single displacement
C. Relationships in the periodic table
1. Periods
2. Groups
3. Representative elements
4. Transition elements
5. Inner transition elements
D. Introduction to organic chemistry
1. Nomenclature of inorganic compounds and coordination compounds
2. Hydrocarbons and functional groups
3. Physical and chemical properties used as examples in empirical / molecular formula calculations,
equilibria, kinetics, colligative properties and stoichiometry
VI. Laboratory experimentation (suggested experimentation)
1. Laboratory safety
2. Determination of the formula of a compound
3. Determination of the percentage of water in a hydrate
4. Determination of molar mass by vapor density
5. Determination of molar mass by freezing point depression
6. Determination of the molar volume of a gas
7. Standardization of a solution using a primary standard
8. Determination of concentration by acid-base titration, including weak acid or weak base.
9. Determination of concentration by oxidation-reduction titration
10. Determination of mass and mole relationship in a chemical reaction
11. Determination of the equilibrium constant for a chemical reaction
12. Determination of appropriate indicators for various acid-base titrations and pH determination
13. Determination of the rate of a reaction and its order
14. Determination of enthalpy change associated with a reaction
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15. Separation and qualitative analysis of cations and anions
16. Synthesis of a coordination compound and its chemical analysis
17. Analytical gravimetric determination
18. Calorimetric or spectrophotometric analysis
19. Separation by chromatography
20. Preparation and properties of buffer solution
21. Determination of electrochemical series
22. Measurement using electrochemical cells and electroplating
23. Synthesis, purification, and analysis of an organic compound.
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