AP Chemistry Syllabus 2013 Mawhiney



Dear AP Chemistry Student and Parent,

Welcome to AP Chemistry!! I am so excited about beginning this journey with you! For those of you that do not know me, my name is Kimberly Mawhiney. This is my 9th year teaching Chemistry and my 3rd year teaching AP Chemistry.

Overview of Course Expectations:

AP Chemistry is a college level course. Students may earn up to 8 hours college credit for successful completion of this course and a good score on the AP Exam. It is a time-consuming and challenging, yet extremely rewarding, course. This course moves at a very fast pace and classroom attendance is a MUST. I will do my very best to provide a college level course/experience which not only prepares you for the AP exam, but provides a solid foundation in chemistry. I also intend for it to be fun!!

The 2014 AP Chemistry Exam will be Monday, May 5 (tentative date)! To be successful on the AP exam, students will need to spend on average five to ten additional hours per week outside of class working on AP Chemistry. This statement is not meant to discourage, but to point out and state the truth to avoid any misconceptions about the high expectations for this course.

Each student will need to have the following supplies:

• # 2 pencils

• Blue/Black Pens

• Composition notebook for labs (We WILL write formal lab reports in this notebook)

• Ear buds for Chrome books

• 3 Ring Binders with LOTS of Paper!

• Scientific Calculator

It's important to recognize that chemistry is a problem-solving class. Major themes and principles are stressed, and one major goal is that the student will be able to apply these principles to understand and solve problems. You should understand that in a science course, a significant portion of your time will be spent solving problems. THIS IS A MATHEMATICS BASED COURSE, and mathematics will be used to solve problems.

AP Chemistry: Six Big Ideas

1) The Chemical Elements are fundamental building materials of matter, and all matter can be understood in terms of arrangements of atoms. These atoms retain their identity in chemical reactions.

2) Chemical and physical properties of materials can be explained by the structure and the arrangement of atoms, ions, or molecules and the forces between them.

3) Changes in matter involve the rearrangement and/or reorganization of atoms and/or the transfer of electrons.

4) Rates of chemical reactions are determined by details of the molecular collisions.

5) The laws of thermodynamics describe the essential role of energy and explain and predict the direction of changes in matter.

6) Any bond or intermolecular attraction that can be formed can be broken. These two processes are in dynamic competition, sensitive to external conditions and external perturbations.

I look forward to getting to know each of you! We will have fun and we will work hard. I check my e-mail frequently, so feel free to email me a question or concern anytime. I also maintain a classroom website that can be accessed from the Currituck County High School home page. Please pay attention to this webpage as it will have information that is pertinent to this class. Assignments, power points, notes, and a timeline of what we will be doing and when is updated weekly on the AP page.

My e-mail address is: kmawhiney@currituck.k12.nc.us

Tentative Schedule: I say tentative, because there have been changes to the AP curriculum and I may be changing the order or time needed for each Unit as I see fit.

1. Chemistry I Fundamentals (3.5 weeks)

Big Ideas: 1, 2, 3, 6

I.  Laboratory Safety

II.  Measurement topics

III.  Symbols and formulas

IV.  Basic Atomic theory & Periodic table

V.  Ionic and covalent bonds

VI.  Nomenclature

VII.  Reactions

VIII. Stoichiometry

       A. Percent composition

       B.  Empirical formulas

       C.  Solutions

       D.  Mole relationships

              1.  % yield

              2.  Limiting reagents

        E.  Titrations and other analyses

IX.  Gases Laws

      A.  Ideal gases

      B.  Boyle’s law

      C.  Charles’ law

      D.  Dalton’s law of partial pressure

      E.  Graham’s law

      F.  Henry’s law

      G.  Van der Waal’s equation of state

X.  Kinetic-Molecular theory

      A.  Avocado’s hypothesis and the mole concept

      B.  Kinetic energy of molecules

      C.  Deviations from ideality

The student will:

1. Define terms such as matter, energy, element, compound, mixture, solution.

2.  Work comfortably with the metric system and work problems using dimensional analysis.

3.  Understand and work with the proper number of significant figures.

4.  Apply knowledge of significant figures to laboratory work.

5.  Correctly use an analytical balance, a vacuum flask, and Buchner funnel.

6.  Know the name and application of the common laboratory equipment used in this course.

7.  Identify the proper safety rules and procedures to be used in experimental situations.

8.  Name the polyatomic ions, given the formula, and vice versa.

9.  Name inorganic compounds, including acids, using the Stock system.

10. Write formulas for the names of inorganic compounds.

11. Work problems involving mole concepts, molarity, percent composition, empirical formulas, and molecular formulas.

12. Balance equations given both reactants and products

13. Solve stoichiometric problems involving percent yield, and limiting reagents.  Apply these concepts to the laboratory setting.

14. State and discuss the major tenants of the kinetic-molecular theory.

15. Apply the kinetic-molecular theory to gases.

16. Discuss the methods and units for measuring pressure; convert between units.

17. Work problems using:  Charles' law, Boyle's law, Gay-Lussac's Law, Avogadro's Law, Dalton's Law, the Ideal Gas Law, and van der Waal's equation.

2. Types of Chemical Reactions and Solution Stoichiometry (3 weeks)

Big Ideas: 1, 3

I.  Reaction types

     A.  Acid base reactions

            1.   Concepts of

                   a) Arrhenius

                   b)  Lowry-Brønsted

      B.  Precipitation reactions

      C.  Oxidation reduction reactions

              1. Oxidation number

              2. Electron transport

D. 5 Basic Types of inorganic reactions

E. Environmental and societal issues involved with reactions

II.  Solution Stoichiometry

III.  Net ionic equations

IV.  Balancing equations including redox

V.   Mass-volume relationships with emphasis on the mole.

The student will:

1.  Apply the periodic law to chemical reactivity in predicting reaction products.

2.  Discuss the activity series of the elements.

4.  Classify compounds as to acids, bases, acid anhydrides, basic anhydrides, salts, and covalent molecules.

5.  Use the properties of metals and nonmetals to predict reaction products.

6.  Write chemical equations for synthesis, decomposition, single replacement, metathetical, redox, combustion, and acid-base reactions.

7.  Use the Periodic Table to predict common oxidation states.

8.  Use the Activity series of elements to predict single replacement reactions.

9.  Know the major components of the atmosphere.

10. Describe physical and chemical properties of reactants and/or products in a reaction.

11. Identify and utilize experimental evidence to determine and describe products of reactions.

12. Understand which ions make water "hard" and know methods of softening water.

3. Chemical Thermodynamics (2.5 weeks)

Big Ideas: 5

I.  State functions

A. Enthalpy

1. Thermal energy, heat, and temperature and work

2. Calorimetry

3. Enthalpy changes

4. Hess’s Law

5. Bond Energies

B. Entropy

II.  Laws of thermodynamics

III.  Gibb’s Free Energy

A. Relationship of enthalpy and entropy to spontaneity of reactions

A. Relationship of change of free energy to equilibrium constants (intro).

B. Relationship of change of free energy to electrode potentials (intro)

The student will:

1. List and define the meanings and common units for the common thermodynamic symbols. Learn the meaning of the following thermodynamic terms:  enthalpy, exothermic, endothermic, system, surroundings, universe, and heat of formation, heat of reaction, calorimetry, heat, calorie, joule, standard molar enthalpy of formation, molar heat of combustion, entropy, absolute entropy, free energy.

2.  Distinguish between a state function and a path function.

3.  Define internal energy, PV work, enthalpy, entropy, and free energy.

4.  Use Hess's law to solve problems of energy, entropy, and free energy.

5.  Define the terms exothermic & endothermic.

6.  Determine the spontaneity of a reaction.

7.  Discuss the laws of thermodynamics (in order).

8.  Solve calorimetry problems involving specific heat.

9. Use stoichiometric principles to solve heat problems.

10. Use enthalpy changes to calculate bond energies.

4. Atomic Structure and Periodicity (3 weeks)

Big Ideas: 1, 2

I.  Electronic Structure

     A. Spectroscopic Evidence for the atomic theory

     B.  Atomic masses

     C.  Atomic number and mass number

     D.  Electron energy levels and orbitals

     E.  Periodic relationships

II.  Nuclear structure

     B.  Half-lives

    

The student will:

1. Identify the major subatomic particles in an atom.

2.  List the types of radioactive emissions.

3.  Discuss the Bohr model of the atom, and compare it to the quantum mechanical model of the atom.

4.  Discuss the major differences in the classical mechanical model and the quantum mechanical model.

5. Investigate the spectroscopic evidence for the modern atomic theory: mass spectroscopy, PES, absorption/emission spectroscopy, IR

6.  Work problems involving energies of electron transitions and apply to ionization energy and PES.

7.  Define and discuss the following terms or concepts:  Heisenberg Uncertainty Principle, Pauli Exclusion Principle, wave-particle duality of matter, Wave function of electrons (Y), radial probability, density, orbitals, aufbau process, and Hund's rule.

8.  Know the shapes of the s, p, and d orbitals.

9.  Understand the basis for the periodic law, and apply it to periodic trends such as atomic radii, ionization energy, electron affinity, melting point, oxidation states, and electronegativity.

10, Work problems involving half-life.

5. Bonding and Molecular Structure (2 weeks)

Big Ideas: 1, 2, 3, 6

I.  Binding forces

      A.  ionic

      B.  covalent

      C.  metallic

      D.  interparticle

II.  Relationships to states, structure, and properties of matter

III.  Polarity of bonds, Electronegativities

IV.  Molecular models

       A. Lewis structures

1. Resonance

       B. Hybridization of orbitals

1. sigma and pi bonds

2. bond order

V.  VSEPR

      A.  Geometry of molecules and ions

      B.  Examples of structural, geometric, isomerism in:

           1. Organic molecules

VI.  Polarity of molecules

 The student will:

1. Draw Lewis structures for the common atoms, ions, and molecules.

2.  Use periodic trends of electronegativity to predict bond type.

3.  Distinguish between polar and nonpolar molecules.

4.  Use electronegativity values and bonding concepts to determine oxidation states on atoms.

5.  Draw resonance structures

6. Use the VSEPR model to predict molecular geometry.

7. Use the hybridization theory to predict molecular geometry.

6. The Kinetic-Molecular Theory and States of Matter (2 weeks)

Big Ideas: 1, 2, 5, 6

I. Interparticle Forces

A. Covalent network

B. Ion-Ion

C. Metallic

D. Van der Waals

II. Relation of molecular structure to physical properties.

III. Liquids and solids

      A.  Liquids and solids from the K-M viewpoint.

    B.  Changes of state.

The student will:

1. State and discuss the major tenants of the kinetic-molecular theory.

2.  Apply the kinetic-molecular theory to liquids and solids.

3.  Discuss intermolecular forces and relate them to physical properties such as boiling point and vapor pressure.

4.  Interpret heating curves as to melting point, boiling point, and specific heat.

5.  Discuss the phenomena of boiling, and be able to relate it to pressure.

6. Explain physical and chemical properties of substance using intra and intermolecular bonding.

Approximate End of 1st Semester

7. Solutions (1.5 weeks)

Big Ideas: 2, 3

I.  Types of solutions

II.  Factors affecting solubility

III. Concentration Expressions

IV. Raoult’s Law

A. Vapor pressure

B. Non-ideal solutions

VI. Colloids

The student will:

1. Define solution vocabulary.

2.  Discuss the effect that physical conditions have on solubility.

3.  Use the concepts of intermolecular forces in discussing the dissolving process and heat of solution.

4.  Separate compounds into electrolytes and non-electrolytes; separate electrolytes into ionic salts, acids, bases, acid anhydrides, and basic anhydrides.

5.  Solve problems involving molarity, % composition, and mole fraction; to be able to convert between concentration designations.

8. Chemical Kinetics (2.5 weeks)

Big Ideas: 1, 3, 4, 6

I.  Rate of reaction

A. Differential Rate Law

B. Integrated Rate Law

II.  Order of the reaction

III. Factors which change the rate of the reaction

        A.  Temperature

        B.  Concentration

        C.  Nature of substance

        D.  Catalysts

IV. Relationship between the rate-determining step and the reaction mechanism

The student will:

1. List the factors that influence the rate of a chemical reaction.

2.  Use experimental data to determine the rate law, determine the order of the reaction, and to define proper units for the constant.

3.  Compare and contrast zero, first, and second order reactions in terms of the plot needed to give a straight line, the relationship of the rate constant to the slope of the straight line, and the half-life of the reaction.

4.  Use experimental data to postulate a reaction mechanism.

5.  Interpret how changing the conditions of the reaction (i.e. temperature, pressure, concentration, and addition of a catalyst) affects both the rate and the rate constant of the reaction.

6.  Discuss the role of a catalyst in the rate and mechanism of a reaction; distinguish between a homogeneous and a heterogeneous catalyst.

7.  Interpret data from a first order reaction to determine its half-life.

9. General Equilibrium (2 weeks)

Big Ideas: 3, 4, 6

I.  Concept of dynamic equilibrium including Le Chatelier’s Principle

II. Equilibrium constants and the law of mass action

A. Kc calculations

B. Kp calculations for gases

C. Reaction Quotient, Q

III. Revisit Free Energy and K

The student will:

1. Describe the meaning of physical and chemical equilibrium, and give real life examples of each.

2.  Write the law of mass action for any system at equilibrium.

3.  Understand the meaning of equilibrium constant and reaction quotient (Q).

4.  Interpret the position of equilibrium from the size of the equilibrium constant.

5.  Use Le Chatelier's Principle to predict the direction a system in equilibrium will shift in order to re-establish equilibrium.

6.  Know that temperature, pressure, and concentration will shift the position of equilibrium.

7.  Understand that a catalyst will not have an effect of the equilibrium constant.

10. Aqueous Equilibria (3.5 weeks)

Big Ideas: 3, 4, 6

I.  Acid Base Theories

    A.  Arrhenius theory

B.  Lowry-Brønsted theory

1. Amphiprotic species

2. Relative strengths of acids and bases

3. Polyprotic acids

II.  Weak Acids and Bases Equilibria

      A.  pH

      B.  pOH

      C.  Buffer systems

      D.  Hydrolysis

E. Titration

III.  Solubility Product Equilibria

      A.  Factors involving dissolution

      B.  Molar solubility

C. Precipitation

The student will:

1. Distinguish between the various modern theories of acids and bases.

2.  Name and write formulas for normal salts, hydrogen salts, hydroxy salts, oxysalts and acids.

3.  Perform a titration and solve for the appropriate concentration.

4.  Use the concept of conjugate acid-base pairs to predict reaction products.

5.  Define and give examples of amphiprotic species.

6.  Identify weak electrolytes.

7.  Write a law of mass action for any reaction in equilibrium.

8.  Know and use the water constant, Kw.

9.  Define pH, pOH, pK, Ka, Kb, ionization constant, percent ionization, Ksp.

10.  Convert from [H3O+] or [OH-] to pH or pOH.

11.  Use a pH meter to determine a titration curve and an ionization constant.

12.  Pick a suitable indicator for a titration.

13.  Recognize salts that undergo hydrolysis and write a reaction for the ion with water.

14.  Given the concentration and amount of weak acids or bases and an appropriate titrant, calculate data to produce a titration curve.

15. Write solubility product expressions for slightly soluble compounds.

16. Solve problems involving: (a) solubility product constants from solubility; (b) molar solubility from Ksp; (c) concentrations of substances necessary to produce a precipitate; (d) concentrations of ions involved in simultaneous equilibrium.

11. Electrochemistry (2 weeks)

(Chapter 17)

Big Ideas: 1, 3, 5, 6

I. Redox equations

II. Galvanic cells and cell potentials

A. Standard Half-Cell Potentials

B. Concentration Cells

C. Free Energy and spontaneity

III. Electrolytic cells

A. Electrolysis (molten and aqueous salts)

B. Corrosion

C. Electroplating and stoichiometric calculations

The student will:

1.  Use the half-reaction method to balance redox equations.

2.  Define electrochemical terms:  redox, anode, anion, cathode, cation, oxidizing agent, reducing agent, emf, and electrode.

3.  Distinguish between an electrolytic cell and a voltaic cell in terms of function, direction and ΔG.

4.  Solve problems using Faraday's law.

5.  Predict reaction products for both electrolytic and voltaic cells.

6.  Use a table of Standard Reduction Potentials to compute cell voltages.

7.  Diagram voltaic cells using proper notation.

8. Establish the relationship between the free energy change, the cell potential, and the equilibrium constant.

Laboratory Experimentation:

Labs form a foundation for student understanding of the chemical principles discussed in lectures but are also chosen to reflect the diversity of lab work generally completed in a first year course. Analysis of data from AP Chemistry examinees shows that increased laboratory time is correlated with higher AP grades. Depending on the particular lab, students will work individually or collaboratively to physically manipulate equipment and materials in order to make relevant observations and collect data. The majority of lab work will be hands-on with simulations used only when the techniques or chemicals cause a hazard. Technology is integrated into a number of labs in the form of probe ware and data collection software.

Student labs are done by lab groups of two to three students. This allows for collaboration and cooperative learning during labs.

The keeping of a lab notebook is required and is designed for the students to present to appropriate staff when enrolled in the college or university of their choice. In this notebook students will communicate lab purpose, safety, procedure, data and observations, calculations, and conclusions. The emphasis in the conclusion will be in three areas. First, students will summarize the lab and determine how and if their purpose was achieved. Second, students will discuss results in terms of their reasonability and how they compare to accepted values (where applicable) and sources of error are identified and discussion of the impact of the error on the final results. Lastly, students identify the major chemical principles used in the lab and the lab results that support those principles.

Although the teacher uses many demonstrations throughout the year, they do not take the place of laboratory work by the students nor are they treated as a lab in the course.

Useful Websites:

There is a multitude of awesome chemistry resources available via the Internet. With the hundreds of tutorial websites out there, I feel confident that you will find adequate information on any topic within this packet. In addition, below are some general chemistry websites for you to check out. You may want to bookmark these and use them as a reference throughout the year.













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