AP Chemistry
2014-2015
Preliminary Syllabis (topics and schedules subject to change)
Text: Chemistry-The Central Science (AP Edition) – 10th Edition
Overview of AP Chemistry
Congratulations! You’ve accepted the challenge offered through a genuine college-level science course. The good
news is that such a course fits your “results-driven” or “challenge-driven” personality on many levels. The good/bad
news is, the course will require much more effort on your part when compared to any other science course you’ve
tackled thus far in high school.
This course is designed to provide a solid first-year (and more) college chemistry experience, both
conceptually and in the laboratory. The labs serve to supplement the learning in the topic section of the course.
Problem-solving skills, both on paper and in the lab, are emphasized. Also, assessments (tests, quizzes, etc.) will
increasingly mirror the questions and testing time allotments used on the actual AP chemistry exam to be taken
next spring. There are on average, one weekly AP chemistry lab completed. During the 4th quarter of the school
year, students will complete several graded practice AP exams, with much less weekly lab-work. The exams are
graded/reviewed in class with AP exam scoring methods so as to increase students’ awareness of test taking
strategies.
Requirements
Students are expected to be in class on time with the required materials. The AP chemistry class meets for 75
minutes 5 days each week. Successful students in the past have recommended a minimum of 45-minutes per day
outside of class set aside for review/homework. Study sessions will be increasingly scheduled outside of class as
the year continues. Students are encouraged to attend as many sessions as necessary.
Laboratory Requirements
Each week, at least one of the class periods (through the 3rd 9 weeks) will be devoted to laboratory work.
Students must keep a formal laboratory notebook (see below-more information in class).
Monthly Bulletin Board Presentation
One time throughout the year, students are responsible for designing and posting visually exciting and
relevant information about a specific topic, which I assign (hint: it will be an AP topic that a student needs more
help with or exposure to).
Students in AP chemistry are expected to take the AP chemistry exam during May 2012.
Laboratory Primer (Remember: Doing chemistry is doing lab…there is no difference)
At the beginning of the year, a laboratory safety guide is provided to students. This safety guide needs to be
reviewed and signed by students and parents/guardians prior participating in any laboratory exercise. Throughout
the laboratory experience, students will keep a formal laboratory notebook (more on this in class). This notebook is
graded with each lab. Laboratory work counts as roughly 25-30% of this class. This laboratory notebook goes with
the student to the university to possibly be used to evaluate his or her placement in a college laboratory program.
Grading
Increasing during the year, grading of all coursework will mimic the style and grading values/conversions used on
AP exams. This style of grading should be fully implemented by the second 9 weeks of school. AP tests, quizzes,
assignments, and bulletin boards may account for approx. 60% of the overall class grade. Lab notebooks,
assignments and practical exams account for approximately 40% of overall class grade. More discussions in class.
Goals of the course (aka “…student will” statements)
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Students will prepare to be critical and independent thinkers who are able to function effectively in a
scientific, technological, and college-based setting.
Students will analyze scientific issues using scientific problem solving
Students will have the resources, capability, and background material necessary to achieve the best
possible personal score on the AP Chemistry examination in May
During AP chemistry laboratory exercises, students will physically manipulate equipment and materials
in order to:
-make relevant observations and collect data
-used the collected data to form conclusions and verify hypotheses
-communicate and compare results and procedures (formally and informally)
Keep a neat and accurate laboratory notebook throughout the year. Students will keep this notebook
beyond high school so as to use as a possible college placement/evaluation tool in the coming years.
Class Schedule*
(*subject to change)
Fall Semester
1st 9 weeks
Unit 1: Introductory and review concepts (2 to 2½ weeks)
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review of summer assignment concepts and AP chemistry prerequisites
Chapters 1-3/REVIEW
Measurement topics, atomic theory, symbols and formulas, periodic table and trends, ionic and covalent bonds,
nomenclature, reactions, stoichiometry: (% composition, complete and net ionic equations, solutions, % yield,
limiting reagents, solubility rules, etc.)
 Expectations discussed, study aids distributed
Unit 1: The student will:
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Define terms such as matter, energy, element, compound, mixture, solution, etc
Work comfortably with the metric system. Work problems using dimensional analysis
Understand and work with the proper numbers of significant figures
Correctly indentify and be able to effectively use laboratory apparatus
Name the polyatomic ions, given the formula, and vice versa
Including the acids, be able to name inorganic compounds from formula, and vice versa
Work problems involving mole concepts, percent composition, empirical formulas, and molecular
formulas
8. Balance equations given both reactants and products
9. Solve stoichiometric problems involving percent yield, and limiting reactants
Apply these concepts to a laboratory setting.
Laboratory:
1. Percent Composition of sugar in bubble gum (review)
2. AP Chemistry Lab 9-Determination of Stoichiometry of Chemical Reactions
1st 9 weeks
Unit 2: Chapter 4 (2 to 2 ½ weeks)
1. Chapter 4-Aqueous Reactions: General properties of aqueous solutions (solubility of ions,
molecules, mixtures), precipitation reactions based on solubility rules, acids, bases, and neutralization.
Students will evaluate oxidation-reduction reactions and predict oxidation numbers. Finally, students will
practice predicting chemical analysis of solutions.
Chapter 4: The student will:
1. Review and experience aqueous solution properties: what/how regarding substances in water (ions,
molecules, mixtures)
2. Identify and predict precipitation reactions (know the solubility rules!)
3. Jump into acids, bases, and neutralization reactions
4. Examine and understand electron-transfer reactions (oxidation-reduction reactions)
5. Correctly perform solution concentration calculations (primarily molality)
6. Experience basic titration concentration methods and lab exercise as ways to complete common
chemical analysis
Apply these concepts to a laboratory setting.
Laboratory or Demonstration:
1. How will electricity travel through various solutions? (Conductivity testing lab)
2. TBA
1st 9 weeks
Unit 3: Chapter 5 (2 to 2 ½ weeks)
Chapter 5-Thermochemistry: First evaluate the nature of energy and the forms it takes (kinetic and
potential). Students will incorporate the terms of work, system, and “surroundings”. The first law of
thermodynamics will be experienced after students practice using terms used in measurements. Enthalpy (heat)
changes will be visualized and calculated in classroom and lab setting. Enthalpies of Reaction describe heat
changes in chemical reactions. Students will construct successful calorimeters to measure and calculate heat
transfer. Enthalpy changes in reactions can be used to establish trends in reaction thermochemistry (using Hess’s
Law). Finally, we will examine foods and fuels as sources of energy.
Chapter 5: The student will:
1. Review the nature of energy and the forms it takes (kinetic and potential), as well as heat systems
2. Explore the first law of thermodynamics (energy cannot be created nor destroyed, only transferred).
Also, the energy possessed by a system will be evaluated.
3. Encounter a state function called enthalpy and Enthalpies of Reaction as a way to quantitatively
measure heat changes in a system. These heat changes predict and reveal a myriad of chemical
characteristics.
4. Examine calorimetry, an experimental technique used to measure heat changes to a chemical process
5. Apply and practice with Hess’s law to predict enthalpy changes of related reactions. Enthalpies of
Formation of compounds will show heat changes in reactions
6. Bring the concepts back to a useful way of looking at foods and fuels
Apply these concepts to a laboratory setting.
Laboratory or Demonstration:
1. General “coffee cup” calorimetry lab
2. AP Chemistry Lab 13: Thermodynamics-Enthalpy of Rxn and Hess’s Law
1st 9 weeks
Unit 4: Chapter 24 (approximately 1 week)
Chapter 24-Organic Chemistry: First evaluate the structure, nature, and nomenclature of organic
compounds (C-H bonded compounds). Students will evaluate compounds with identical compositions but different
structures (isomers). Many organic compounds contain functional groups, a group of atoms at which most of the
compound’s chemical reactions occur. Students can visualize that compounds with nonsuperimposable mirror
images are chiral. Chirality plays an important role in organic and biological chemistry. Students will experience
biological chemistry as they evaluate proteins, carbohydrates, lipids, and nucleic acids. This extensive branch of
chemical science will be further evaluated using research
Chapter 24: The student will:
1. Evaluate the structure, nature, and rules for nomenclature of organic compounds (C-H bonded
compounds)
2. Briefly evaluate compounds with identical compositions but different structures (isomers). Note: many
isomers with identical compositions retain entirely different chemical properties
3. Briefly experience the nuances and application of functional group chemistry, a functional group is a
group of atoms at which most of the compound’s chemical reactions occur.
4. Visualize that compounds with non-superimposable mirror images are chiral. Chirality plays an
important role in organic and biological chemistry.
5. Evaluate and research the billion(s) -dollar branch of chemistry called biological chemistry, including
the nature of organic molecules of proteins, carbohydrates, and lipids
Apply these concepts to a laboratory or research setting.
Laboratory/Demonstration/Research:
1. Functional Group chemistry research project
2. Pool water chemistry lab (as time allows)
End of 1st 9 weeks Exam:
Chapters 1-3: (review concepts from “Chem I” and stoichiometry)
Chapter 4: Aqueous Reactions
Chapter 5: Thermochemistry
Chapter 24: Organic Chemistry (brief)
Note: A laboratory “practical” exam may be given, and count towards 30-40% of 9-weeks’ exam
2nd 9 weeks
Unit 1: Chapter 10 (2 weeks)
Chapter 10-Gases: Students begin by qualitatively distinguishing between the characteristics of gases with
liquids and solids. We then study gas pressure, measurements, and units. The remainder of the unit focuses on the
understanding, functionality, and application of the gas laws in chemistry understanding. Students will perform
labs and simulations that vary the amounts of volume, pressure, temperature, and quantity of matter.
Most gases come close to following the Ideal Gas Law equation: PV=nRT. This can be used to calculate the density
or molar mass of a gas.
In gas mixtures, each gas exerts a pressure (partial pressure) that is part of the total pressure
The kinetic-molecular theory of gases helps us to understand and quantify gas behavior on the molecular level.
Two associated properties are effusion and diffusion
Finally, real gases deviate from ideal gas behavior in that A) gas molecules have finite volume, and B) attractive
forces do exist between gaseous molecules. We quantify this attraction using the van der Waals equation.
This will be a lab-intensive unit
Chapter 10: The student will:
1. Visually compare the general physical and chemical characteristics of solids, liquids, and gases
2. Experience gas pressure, measurements, and unit methodology. The remainder of the unit focuses on
the understanding, functionality, and application of the gas laws in chemistry understanding
3. Perform labs and simulations that vary the amounts of volume, pressure, temperature, and quantity of
matter as an introduction to several important gas laws. Following several simulations, students will
solve for several unknown values using the gas laws, and then be asked to evaluate and explain their
answers
4. Use their lab results in conjunction with background knowledge to describe examples where deviations
from ideal gas law behaviors can be quantified, if possible
Apply these concepts to a laboratory setting.
Laboratory or Demonstration:
1. AP Chemistry Lab 3-Determination of Molar mass of Gases, Liquids
2. AP Chemistry Lab 5-Determination of Molar Mass of a Gas
2nd 9 weeks
Unit 2: Chapter 14 (1 ½ to 2 weeks)
Chapter 14-Chemical Kinetics: Students begin by exploring the variables that affect reaction rates:
concentration
physical states of reactants
temperature
and the presence of catalysts
These factors can be understood in terms of the likelihood of collisions of reactant substances
Students will consider how to express reaction rates and reaction stoichiometry
The effect of concentration on rate is expressed quantitatively using rate laws
Rate equations can change with time as concentrations change with time by evaluating zero-order, first-order, and
second-order reactions
The usually temperature-related activation energy relates to the minimal input of energy required to initiate a
reaction.
Reaction mechanisms represent the step-by-step molecular pathways leading from reactants to products
In regards to reaction kinetics, catalysts refer to substances that can increase reaction rates. An example of a
biological catalyst is an enzyme.
Chapter 14: The student will:
1. Quantitatively and qualitatively explore the variables that affect reaction rates. These are italicized
above
2. Come to realize through several means that reaction rate outputs are the result of reacting substance
collision rate inputs. Several factors can affect these collision rates.
3. Consider and understand how to express reaction rates and reaction stoichiometry, the effect of
concentration on rate is expressed quantitatively using rate laws
4. Through collaboration and evaluation in both classroom and laboratory setting, discover that rate
equations can change with time as concentrations change with time by evaluating zero-order, firstorder, and second-order reactions. Furthermore, the usually temperature-related activation energy
relates to the minimal input of energy required to initiate a reaction.
5. Through collaboration and evaluation in both classroom and laboratory setting, discover the reaction
mechanism steps of a reaction, as well as the application of the related catalysts and enzymes to the
reaction process
Apply these concepts to a laboratory setting.
Laboratory or Demonstration:
1. AP Chemistry Lab 12 - Kinetics of a Reaction
2nd 9 weeks
Unit 3: Chapter 15 (2 to 2 ½ weeks)
Chapter 15-Chemical Equilibrium: Students begin by examining reversible reactions and the concept of
equilibrium.
Through coursework and laboratory work, we examine the equilibrium constant based on forward and reverse
reactions for both homogeneous and heterogeneous reactions
The next step is the writing of equilibrium constant expressions so as to complete quantitative analysis in
chemistry
Applications of equilibrium constant chemistry will be explored. Equilibrium constants can used to predict
equilibrium concentrations of reactants and products, as well as reaction direction.
Le Chatelier’s principle predicts how a system at equilibrium responds to changes in concentration, pressure,
volume, and temperature
Chapter 15: The student will:
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Describe the meaning of physical and chemical equilibrium, and give real life examples of each
Write the law of mass action for any system at equilibrium
Understand the meaning of equilibrium constant and reaction quotient (Q)
Interpret the position of equilibrium from the size of the equilibrium constant
Use Le Chatelier’s principle to predict the direction a system in equilibrium will shift in order to reestablish equilibrium
6. Know that temperature, pressure, and concentration will shift the position of equilibrium
7. Understand that a catalyst will not have an effect on the equilibrium constant.
Apply these concepts to a laboratory setting.
Laboratory or Demonstration:
1. AP Chemistry Lab 17 – Determination of Keq for FeSCN
2nd 9 weeks
Unit 4: Chapter 16 (2 to 2 1/2 weeks)
Chapter 16-Acid-base Equilibria: Students begin by examining the Arrhenius definition of acids and
bases
Through coursework and laboratory work, we learn that a Bronsted-Lowry (B-L)acid is a proton donor, and that a BL base is a proton acceptor. These chemical species differ as conjugate acid-base pairs. We use the pH scale to
describe the acidity or basicity of an aqueous solution. Neutral solutions have a pH =7. Strong acids and bases are
strong electrolytes, ionizing or dissociating completely. Weak acids and bases are weak electrolytes, and ionize
only partially. We learn that the ionization of weak acids and bases in water is an equilibrium process with
equilibrium constants (Ka or Kb) that can be used to calculate the pH of a solution. We see that the relationship Ka
x Kb =Kw means that the stronger an acid, the weaker it’s conjugate base. The relationship between chemical
structure and acid-base behavior is explored. Finally, we learn that a Lewis acid is an electron-pair acceptor, and a
Lewis base is an electron-pair donor.
Chapter 16: The student will:
1. Distinguish between the various modern descriptions of acids and bases
2. Name and write formulas for various salts and acids
3. Write a law of mass action for any reaction in equilibrium
4. Write balanced equations involving acids, bases, and salts
5. Define pH, pOH, pK, Ka, Kb, ionization constant, percent ionization, Ksp
6. Convert from (H3O+) or (OH-) to pH or pOH
7. Use the concept of conjugate acid-base pairs to predict reaction products
8. Define and give examples of amphiprotic species
9. List the six strong acids, and know what defines them as “strong”
10. Recognize Lewis acid-base reactions
Apply these concepts to a laboratory setting.
Laboratory or Demonstration:
1. AP Chemistry Lab 12 - Determination of an Ionization Constant (Ka) of weak acids
2. AP Chemistry Lab 6 – Acid-Base Titrations
End of 2nd 9 weeks Exam:
1. Chapters 1-3: (review concepts from “Chem I” and stoichiometry)
2. Chapter 4: Aqueous Reactions
3. Chapter 5: Thermochemistry
4. Chapter 24: Organic Chemistry
5. Chapter 10: Gases
6. Chapter 14: Kinetics
7. Chapter 15: Equilibrium
8. Chapter 16: Acid/Base Equilibrium
Note: A laboratory “practical” exam may be given, and count towards 40% of semester exam
3rd 9 weeks
Unit 1: Chapter 17 (2 to 2 ½ weeks)
Chapter 17-Additional Aqueous Equilibrium: We begin by considering a specific example of Le
Chatelier’s principle known as the common-ion effect.
We consider buffered solutions and explore how they resist pH change when small amounts of a strong acid or
base are added.
We examine acid-base titrations and explore how to determine pH at any point in an acid-base titration
We learn how to use solubility-product constants to determine to what extent a sparingly soluble salt dissolves in
water.
We investigate some of the factors that affect solubility, including the common-ion effect and the effect of acids
We learn how differences in solubility can be used to separate ions through selective precipitation
Finally, we explain how the principles of solubility and complexation equilibria can be used to identify ions in
solution
Chapter 17: The student will:
1. Examine the common-ion effect in chemical reactions
2. Learn how buffered solutions resist pH changes even when small amounts of a strong acid or base are
added.
3. Perform acid-base titrations and explore how to determine pH and equivalence point
4. Examine how to use solubility-product constants to determine salt dissolution in water
5. Investigate some of the factors that affect solubility
6. Learn that solubility differences can be used to separate ions in solution and identify ions
Apply these concepts to a laboratory setting.
Laboratory or Demonstration:
1. AP Chemistry Lab 11 - Selecting Indicators for Acid-Base Titrations
3rd 9 weeks
Unit 2: Chapter 19 (2 weeks)
Chapter 19 – Chemical Thermodynamics: Changes/reactions in nature often have a directional nature.
Some reactions are spontaneous in a particular direction. Studies in entropy and the second law of
thermodynamics shows that entropy of the universe increases. The third law of thermodynamics states that at 0
Kelvin, the entropy of a crystalline solid is zero. Entropy values for system changes can be calculated. Gibbs free
energy is a useful value for determining how far a chemical system is from equilibrium. Finally, we consider how
the standard free- energy change for a reaction can be used to calculate the equilibrium constant for the reaction.
Chapter 19: The student will:
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List and define the meanings and common units for the common thermodynamic symbols
Evaluate internal energy, PV work, enthalpy, entropy, and free energy
Use Hess’s law to solve problems of energy, entropy, and free energy
Re-examine the terms exothermic, endothermic, exergonic, and endergonic
Determine the spontaneity of a reaction
Discuss the laws of thermodynamics
Evaluate the relationship between free energy change and equilibrium constants
Apply these concepts to a laboratory setting.
Laboratory or Demonstration:
3rd 9 weeks
Unit 3: Electrochemistry (2 to 2 1/2 weeks)
Chapter 20 – Electrochemistry: Oxidation states of substances, as well as oxidation-reduction reactions
(redox) must first be reviewed. The half-reactions method of balancing redox reactions can then be used. Voltaic
cells (or batteries) produce electricity from spontaneous redox reactions. The electrode where oxidation (lose
electrons) occurs is called the anode, while the electrode where reduction (gain electrons) takes place is called the
cathode. Cell potential (difference of electrical potentials at the two electrodes) is a measure of volts. We
reference standard reduction potentials to calculate cell potentials under standard conditions (under non-standard
conditions, we can examine cell potential using the Nernst equation. We relate Gibbs free energy to cell potential.
In addition, battery and fuel cell technology are examined as important energy sources that use
electrochemical reactions. Corrosion is a common example of redox electrochemistry.
Chapter 20: The student will:
1. Review oxidation number calculations as an important part of redox chemistry
2. Use the half-reaction method (as well as the arrow method) as a tool to help balance redox eqtns
3. Define electrochemical terms: redox, anode, anion, cathode, cation, oxidizing agent, reducing agent,
emf, electrode, etc.
4. Distinguish between an electrolytic cell and a voltaic cell in terms of G
5. Predict reaction products for electrolytic and voltaic cells.
6. Solve electrolytic problems using Faraday’s law and the Nernst equation.
7. Use a table of Standard Reduction Potentials to compute cell voltages
8. Diagram voltaic cells using proper notation
9. Establish the general relationship between the free energy change, the cell potential, and the
equilibrium constant.
Apply these concepts to a laboratory setting.
Laboratory or Demonstration:
1. AP Chemistry Lab 8 – Oxidation-Reduction Titrations
2. Simple voltaic experiment(s)
3rd 9 weeks
Unit 4: Aqueous Solutions (1 week)
Chapter 13 – Solutions: We will look at what happens at the molecular level when substances dissolve,
paying particular attention to the role of molecular forces. The amount of solute in a saturated solution defines the
solubility of the solute, the extent to which a particular solute dissolves in a particular solvent. Many factors can
affect solubility, including pressure, temperature, and chemical phase. Can be expressed in several different ways,
including mole fraction, molarity, and molality. Some physical properties of solutions depend upon the
concentration of the solute, it’s identity. These colligative properties include the extent to which the solute lowers
the vapor pressure, increases the boiling point, and decreases the freezing point of the solvent. Finally, colloids are
mixtures that are not true solutions, but rather consist of at least one phase of particles larger than typical
molecular sizes.
Chapter 13: The students will:
1. Examine what happens on the molecular level when substances dissolve.
2. Perform qualitative and quantitative evaluations of the role of intermolecular forces on solutions, as
well as a review of the types and strengths of these forces
3. Work with solutions and stoichiometry to learn about solutions and equilibrium, including the factors
that affect both
4. Examine several common ways of expressing concentration, including mole fraction, molarity, and
molality
5. Perform laboratory(s) and calculations to show the relationships between solution concentration and
colligative properties.
6. Close the chapter by briefly examining the properties of colloids.
Apply these concepts to a laboratory setting.
Laboratory or Demonstration:
1. AP Chemistry Lab Challenge - Who can create the most “effective” antifreeze/coolant (colligative
solution)?
2. AP Chemistry Lab 4 – Molar Mass by Freezing Point Depression
End of 3rd 9 weeks Exam:
1. Chapters 17: Additional equilibrium
2. Chapter 19: Thermodynamics
3. Chapter 20: Electrochemistry
4. Chapter 13: Solutions
Last Nine Weeks
Last 9 Weeks-Unit 1: Chapter 11 (2 to 2 ½ weeks)
Chapter 11-Intermolecular Forces: A molecular comparison of solids, liquids, and gases reveals the
important roles of temperature and intermolecular forces plays in determination of physical states of a substance.
We then examine the four primary intermolecular forces:
Dispersion forces
Dipole-dipole forces
Hydrogen bonds
Ion-dipole forces
The nature and strength of the intermolecular forces between molecules are largely responsible for many liquid
properties, including viscosity and surface tension. A dynamic equilibrium exists between a liquid and it’s gaseous
state. The affects of vapor pressure will also be examined. Phase diagrams are graphic representations of the
equilibria between the gaseous, liquid, and solid phases. Some substances pass into a liquid crystalline phase,
which is an intermediate phase between the solid and liquid phases.
Chapter 11: The students will:
1. Examine solids, liquids, and gas from a molecular perspective.
2. Experiment through laboratory observations and other coursework the four primary intermolecular
forces listed above. Central to an understanding of these forces is a deeper relationship of chemical
attraction, repulsion, and energy
3. Then examine the role of intermolecular forces on physical liquid properties such as viscosity and
surface tension
4. Explore the dynamic equilibrium that exists between a liquid and it’s gaseous state. This relationship
will allow students to further explore the concept of vapor pressure.
5. Become familiar with how to read, interpret, and perform later calculations involving phase diagrams
6. Be exposed to the nature and use of liquid crystal technology in today’s society.
Apply these concepts to a laboratory setting.
Laboratory or Demonstration:
1. Odyssey Virtual Lab – Intermolecular Forces
2. Various in-class demonstration(s) of classroom topics
Last 9 Weeks-Unit 2: Chapter 6 (2 weeks)
Chapter 6-Electronic Structure: A brief introduction explains that light (radiant energy or
electromagnetic radiation) has wavelike properties characterized by wavelength, frequency, and speed. However,
this radiation also has particle-like properties and can be described in terms of photons, “particles” of light. Line
spectra indicates that electrons exist only at certain energy levels around the nucleus, and energy is involved when
an electron “moves” from one level to another at fixed positions (Bohr model). Heisenberg’s uncertainty principle
establishes that it is impossible to determine simultaneously the exact position and motion of an electron in an
atom. The wave functions that mathematically describe the electron’s position and energy in an atom are called
atomic orbitals. These can be described using atomic numbers (n). These electron orbitals can be graphically
represented in many different ways, and each orbital can hold a maximum of two electrons.
Knowing orbital energies as well as some fundamental characteristics of electrons described by Hund’s
rule allows us to determine how electrons are distributed in an atom (electron configurations). A direct correlation
exists between the electron configuration of an atom, and it’s location on the periodic table.
Chapter 6: The students will:
1. Discover what is meant when we say that radiant (light) energy can be described by both wave-like and
particle-like properties
2. Discuss the Bohr model of the atom, and compare it to the quantum mechanical model of the atom
3. Work problems involving quantum numbers, and energies of electron transitions
4. Relate the quantum numbers to the number and types of orbitals and recognize the different orbital
shapes
5. Draw an energy-level diagram for the orbitals in a many-electron atom and describe how electrons
populate the orbitals in the ground state of an atom, using the Pauli exclusion principle and Hund’s rule
6. Use the periodic table to write complete and/or condensed electron configurations, and determine the
number of unpaired electrons in an atom.
Apply these concepts to a laboratory setting.
Laboratory or Demonstration:
1. Odyssey Virtual Lab – Orbitals, energy levels, and molecular shapes
Last 9 Weeks
Unit 3: Chapter 7 (1-2 weeks)
Chapter 7-Periodic Properties of Elements: First and foremost, many measureable and visual
trends can be established with elements in the periodic table. These will be evaluated (quantitatively and
qualitatively) throughout the chapter.
Many properties of atoms depend on the net attraction of the outer electrons in the nucleus and on the
average distance from the nucleus. The effective nuclear charge represents the net positive charge of the nucleus.
Ionization energy represents the energy required to remove one or more electrons from an atom. The periodic
trends in ionization energy depend on variations in effective nuclear charge and atomic radii. Elements with higher
ionization energies generally have higher electron affinities as well.
The physical and chemical properties of metals are different from those of nonmetals. These properties
arise from the fundamental characteristics of atoms, particularly ionization energy. Metalloids display properties
that are intermediate between those of metals and non-metals. Trends in individual periodic groups (ex. Group 1A,
2A, etc.) will also be examined.
Chapter 7: The students will:
1. Understand the meaning of effective nuclear charge (Zeff) and how Zeff depends on nuclear charge and
electron configuration
2. Use the periodic table to predict trends in atomic radii, ionization energy, and electron affinity
3. Explain how (in terms of attraction, repulsion, and energy) the radius of an atom changes upon losing
electrons to form a cation, or gaining electrons to form an anion.
4. Be able to write electron configurations of atoms
5. Explain and recognize (quantitatively and graphically) how the ionization energy changes as we remove
successive electrons. Recognize the jump in ionization energy that occurs when the ionization
corresponds to removing a core electron
6. Understand how irregularities in the periodic trends for electron affinity can be related to electron
configuration
7. Recognize the differences in chemical and physical properties of metals and non-metals, including the
basicity of metal oxides and the acidity of nonmetal oxides
8. Understand how the atomic properties are related to the chemical reactivity and physical properties of
the alkali and alkaline earth metals (Group 1A and 2A)
9. Be able to write balanced equations for the reactions of the group 1A and 2A metals with water, O, H,
and the halogens
10. Understand, recognize, and explain the unique characteristics of hydrogen
11. Understand how the atomic properties are related to the chemical reactivity and physical properties of
the group 6A, 7A, and 8A elements
Apply these concepts to a laboratory setting.
Laboratory or Demonstration:
1. Simple flame tests
2. Periodicity and chemistry factors
Last 9 Weeks
Unit 4: Chapter 8 (1.5 to 2 weeks)
Chapter 8-Chemical Bonding: Ionic, covalent, and metallic bonding will be looked at prior to practice with
Lewis structures. Lewis structures are a simple yet powerful way of predicting covalent bonding patterns in
molecules. In addition to the octet rule, we see that the concept of formal charge can be used to identify the
dominant Lewis structure. When more than one equivalent Lewis structure can be drawn for a molecule or ion, the
bonding structure is depicted as a blend of two or more resonance structures.
Ionic Bonding: When learning about ionic compounds, an understanding of electrostatic attractions between ions
allows students to better understand the concept of lattice energy.
Covalent Bonding: Bonding in molecular substances is accomplished by the sharing of one or more electron pairs.
First and foremost, many measureable and visual trends can be established with elements in the periodic table.
These will be evaluated (quantitatively and qualitatively) throughout the chapter. The term electronegativity is
related to the ability of an element in a compound to attract electrons to itself. When electron pairs in a compound
are shared unequally between atoms with different electronegativities, the molecular bond is referred to as polar
covalent bonds.
We will recognize that the octet rule is more of a guideline than an absolute rule. Exceptions to the octet
rule include some molecules where:
A) There is an odd number of electrons
B) Large differences in electronegativities prevent an atom from completing it’s octet, and/or
C) An element from period 3 or below in the periodic table attains more than anoctet of electrons
Bond strengths in covalent compounds vary with the number of shared electron pairs, as well as other
factors. We use average bond enthalpy values to estimate the enthalpies of reactions in cases where
thermodynamic data are unavailable.
Chapter 8: The students will:
1. Write Lewis symbols for atoms and ions
2. Understand lattice energy and be able to arrange compounds in order of increasing lattice energy
based on the charges and sizes of the ions involved
3. Use atomic electron configurations and the octet rule to write Lewis structures for molecules to
determine their electron distribution
4. Use electronegativity differences (and molecular symmetry) to identify nonpolar covalent, polar
covalent, and ionic bonds. Students will also be able to identify coordinate covalent bonds.
5. Calculate charge separation in diatomic molecules based on the experimentally measured dipole
moment and bond distance
6. Calculate formal charges from Lewis structures and and use those formal charges to identify the
dominant Lewis structure for a molecule or ion
7. Recognize where resonance structures are needed to describe the bonding
8. Recognize and be able to describe exceptions to the octet rule and draw accurate Lewis structures
even when the octet rule is not obeyed
9. Understand the relationship between bond type (single, double, and triple), bond strength (bond
enthalpy), and bond length
10. Use bond enthalpies to estimate enthalpy changes for reactions involving gas-phase reactants and
products
Apply these concepts to a laboratory setting. Laboratory or Demonstration:
1. Odyssey Virtual Lab – Lewis structures
2. In-class hands-on working with atomic orbital and crystal shape models
Last 9 Weeks
Unit 5: Chapter 9 (2 weeks)
Chapter 9 - Molecular Geometry: Many molecular shapes are common in nature. We consider how
molecular geometries can be predicted using the valence-shell-electron-pair repulsion (or VSEPR) model, which is
based on Lewis structures and the repulsions between regions of high electron density. Bond type and molecular
geometry allow us to predict whether a molecule is polar or non-polar. In valence-bond theory, the bonding
electrons are visualized as originating in atomic orbitals on two atoms. A covalent bond is formed when these
orbitals overlap.
When atomic orbitals of one atom in a compound “mix” with another, or hydridize, they create hybrid
orbitals.
Atomic orbitals that contribute to covalent bonding in a molecule can overlap in multiple ways to produce
sigma and pi bonds between atoms. Single bonds generally consist of one sigma bond, while multiple bonds
involve one sigma and one or more pi bonds. We examine the geometric arrangements of these bonds.
We will briefly examine molecular orbital theory, which introduces the concepts of bonding and antibonding molecular orbitals.
Chapter 9: The students will:
1. Be able to describe the three-dimensional shapes of molecules using the VSEPR model
2. Determine whether a molecule is polar or nonpolar based on it’s geometry and the individual bond
dipole moments
3. Be able to explain the role of orbital overlap in the formation of covalent bonds
4. Be able to specify the hybridization state of atoms in molecules based on observed molecular
structures
5. Have the ability to sketch how orbitals overlap to form sigma and pi bonds
6. Be able to draw basic molecular orbital energy-level diagrams and place electrons into them to obtain
the bond orders and electron configurations of diatomic molecules using molecular orbital theory
7. Be able to understand and fully explain the relationships among bond order, bond strength (bond
enthalpy), and bond length
Apply these concepts to a laboratory setting.
Laboratory or Demonstration:
1. Odyssey Virtual Lab – Atomic Orbitals
End of Year Exams:
*AP Chemistry Exam through College Board – Date to be Announced
*After-class study sessions to be initiated during 4th quarter
*Laboratory “practical” exam may be given, and count towards approx. 40% of semester grade