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Syllabus for AP Chemistry Teacher: Mr. Lambert Room: M101 Textbooks: General Chemistry 4th Edition. John W. Hill, Ralph H. Petrucci, Terry W. McCreary, Scott S. Perry. Pearson Prentice Hall, 2005. ISBN: 0-13-140283-8 Laboratory: Labs are used or adapted from the following: Laboratory Experiments for Chemistry: The Central Science. John H. Nelson and Kenneth C. Kemp. Prentice Hall Publishing, 2000. ISBN: 0-13-084101-3 Laboratory Investigations: AP Chemistry. David Hostage and Martin Fossett. Peoples Education, 2006. ISBN: 14138-0489-6 Objective: The objective of this course is to build on the foundation of chemical knowledge developed in the first year course by strengthening the student’s basic understanding of chemical concepts, such as atomic structure, gases, chemical kinetics, chemical equilibrium, and thermodynamics. This course aims to: 1) provide the students with a basic understanding of chemical concepts in a manner that is equivalent to a first-year college chemistry course; 2) provide students with a safe and enjoyable learning environment where they feel free to express their ideas, ask questions, and offer answers; 3) strengthen student’s quantitative reasoning and problem-solving skills; 4) promote team-building skills through collaborative laboratory exercises; 5) teach students how to work quickly and effectively when presented with a difficult problem; 6) help students express their answers to theoretical questions clearly and concisely; and 7) develop independent thinkers that are curious about the scientific world around them. Method of Instruction: The majority of the information is presented in a lecture/discussion setting. Throughout the class period, questions and problems will be presented to the student directed towards strengthening their understanding of the topic of discussion. The solutions to these questions and problems are discussed using as much student input as possible. Laboratory experiments are performed approximately once a week during a 60 minute block. At this time, student will: 1) directly examine the concepts discussed in class; 2) learn and use the fundamental skills required to perform chemical experiments; 3) gather and analyze data collected during experiments; and 4) present the results of the experiment in a clear, concise manner. The procedure, any data a student collects, and the analysis of that data for a given lab will be kept in a formal laboratory notebook. Method of Evaluation: Your grade in this class will determined through your performances on unit tests, weekly quizzes, and homework. Every student who takes the class is encouraged to take the SAT II Chemistry subject test, but they are required to take the AP exam. a. Tests: Tests will make up 55% of the semester grade. All tests will be weighted equally. There will be a test at the completion of each unit of study. Each test will be broken down into two sections: multiple choice and free response. Each section will have a time limit. You will start with the multiple choice section. At the end of the allotted time, I will collect the multiple choice section and hand out the free response section. Calculators may only be used on the free response section. b. Quizzes: Quizzes will make up 25% of the semester grade. You will encounter three types of quizzes in this class: in-class, take-home, and lab. In-class quizzes will be closed-book, and most will be multiple choice. The take-home quizzes will be adapted from past AP free response questions dealing with the current topic of discussion. These quizzes will be open-book, open –note and will be collected at the start of the next class period. Lab quizzes are designed to evaluate the student’s understanding of the procedure, data manipulation, and results analysis. The quizzes will also be take-home and collected at the start of the next class period. If you are not finished with your take-home quiz, you may ask to leave class to finish it. If you choose to do this, there will be an automatic 5 percentage point penalty on the quiz. If you do not hand in the quiz by the end of the class period, there will be an addition 5 percentage point penalty AND you will be reported as having an UNEXCUSED absence from class. c. Problem Sets: Problem sets grades will make up the final 20% of the semester grade. Selected questions from the problem set will be due the day after the completion of a section within each unit. These questions will be graded for correctness. In order to earn full credit on the problem set, you must show all your work for every question. THERE ARE NO ONE-WORD ANSWERS IN CHEMISTRY, even though the question may appear to be asking for one. Always explain your answers. If you are having problems, you’ll need to see me outside of class. d. Extra Credit: Extra Credit can be earned for each unit test. This is an optional exercise although it is one that I highly recommend. The amount of extra credit given will vary from test to test, but it will be at least 5 points. Class Expectations: 1. You should always be prepared for class and ready to begin working as soon as the bell rings. We have a lot of material to cover, and I always look forward to seeing your smiling faces. If the door is closed upon your arrival to class, it means that you are tardy. If this occurs more than 5 times over the course of the year, I will double your minutes when I report them to the Dean’s office. Message: Don’t be late! 2. You should keep all notes, handouts, worksheets, homework assignments, tests, and quizzes together in one location, preferably in a 3-ring binder that should be specific for this class. Since this is an AP course, the expectation is that you will master all the material that is covered in the course. In order to master this material, you WILL have to review it, which is why all your material should be located in one place. 3. You are required to bring a calculator to class. We will be working through a lot of problems in class. Failure to bring a calculator to class will result in tardy minutes as I will make you go and get your calculator. 4. You are expected to read your textbook, even though you may not be given specific reading assignments. Your textbook is a great resource, use it. 5. You will treat everyone in the class with respect. This is a hard class, and I know that there will be a lot of questions. Everyone should feel free to ask those questions and participate in class. If, at any time, I feel you are being disrespectful to me or your classmates through your words or actions, I will ask you to leave class (and you will have to deal with the consequences of being dismissed from class). 6. The laboratory is a serious place, and I take it seriously. No horseplay will be tolerated. I expect you to wear you lab goggles at all times. If labs are not completed in the allotted class time, you will be expected to finish them outside of class. 7. Major assignments, such as lab reports and projects, will drop 1 letter grade everyday they are overdue. If you know you will have a problem making a deadline, please talk to me BEFORE the assignment is due. 8. A test or a quiz may be postponed 1 day if you arrive back on campus after 8:30 p.m. from a SCHOOL TRIP the night before the test is to be taken or the quiz is due. I will expect you to make up the test the following day either in the morning before classes or during a free period (take-home quizzes will be due the following day at the start of class). 9. You are expected to make up a missed quiz or test (due to an excused absence) at some point the following day. It is YOUR responsibility to set up a time with me to make up the missed work. Failure to do so will result in a zero in the grade book. In the case of unexcused absence, you will receive a 0 on any work due that day, including any tests or quizzes taken during the class period. 10. You are expected to follow all the school’s policies set forth by the Honor Council and in the Student Handbook. 11. I will allow food and drink into the classroom if it does not disrupt the class. Absolutely no food or drinks will be allowed in the laboratory area. The classroom is not a garbage can. Clean up after yourself. This is a privilege, not a right. If you abuse it, it will be taken away. 12. I’m happy to give extra help. I will be in my classroom every morning by 7:15 am. Come see me! Course Outline: Unit 1: Gases Chapter 5: Gases The student will: 1. Describe the physical properties of a gas. (Sect. 5.1-5.2) 2. Define pressure and convert between various units. (Sect. 5.3) Prob: 21 b & c, 22 a 3. Understand how open and closed-ended manometers work. (Sect. 5.3) Prob: 26, 28 4. Use Boyle’s law to relate pressure and volume to a fixed quantity of gas at a fixed temperature. (Sect. 5.4) Prob: 32 5. Use Charles’s law to relate temperature and volume of a fixed amount of gas at a fixed pressure. (Sect. 5.5) Prob: 38, 39 6. State Avogardo’s law, define STP, and use the molar volume of gas at STP. (Sect. 5.6) Prob: 44, 46, 99 7. List the principles of the kinetic-molecular theory of gases and describe their relationship to the molecular properties of gases and to the simple gas laws. (Sect. 5.2-5.11) Prob: 91 8. Calculate initial and final values of gas variables (n, P, V, T). (Sect. 5.7) Prob: 47, 50 9. Calculate unknown properties of a gas (n, P, V, T) using the ideal gas law. (Sect. 5.8) Prob: 56, 58 10. Be able to manipulate the ideal gas law to calculate the molar mass and the density of a gas. (Sect. 5.8) Prob: 62, 68, 70, 101, 119 (give molecular formula only) 11. Perform stoichiometric calculations involving gases as reactants or products. (Sect. 5.9) Prob: 71, 75, 76, 100, 111 12. Apply Dalton’s law of partial pressures to determine the total pressure of a sample of gas or the partial pressures of individual gases in a mixture of gases. (Sect. 5.10) Prob: 85, 87 13. Calculate the mole fraction of components of a gas mixture. (Sect. 5.10) Prob: 84 14. State Graham’s law and apply it to calculating the relative rates of effusion of different gases. (Sect. 5.11) Prob: 90, 93 a, 120 15. Explain why gases may exhibit nonideal behavior. (Sect. 5.12) 16. Describe the changes that must be made to the ideal gas law when dealing with real gases. (Sect. 5.12) Laboratories to be performed: 1. Determination of the molar mass of a vapor. Unit 2: Solution Chemistry Chapter 12: Physical Properties of Solutions The student will 1. Recognize the three types of solutions that can be formed from the three states of matter. (Sect. 12.1) 2. Describe the influence of intermolecular forces on the energies of solution formation. (Sect. 12.3) Prob: 43, 82 3. Distinguish between ideal and non-ideal solutions. (Sect. 12.3) 4. Predict if a mixture is likely to be homogeneous or heterogeneous based on information about intermolecular forces. (Sect. 12.3) 5. Describe the process by which ionic compounds dissolve in water. (Sect. 12.3) 6. Define solubility; distinguish between saturated, supersaturated, and unsaturated solutions; and extract solubility data from solubility curves. (Sect. 12.4) Prob: 46, 48 7. Describe the effects of temperature and pressure on the solubility of gases. (Sect. 12.5) 8. Calculate the solubility of gases using Henry’s law. (Sect. 12.5) Prob: 51 9. Express solution concentrations in molarity, molality, mole fraction, mass percent, and perform conversions between the different concentration units. (Sect. 12.2) Prob: 28, 35, 36, 38 Chapter 4: Chemical Reactions in Aqueous Solutions Problems: 25, 26, 29, 33, 35, 42, 46, 47, 56, 60, 63, 68, 70, 68, 76, 77, 83, 92, 102, 113, From Ch. 18: 27, 29 The student will 10. Identify strong electrolytes, weak electrolytes, and nonelectrolytes. (Sect. 4.1) Prob: 25 11. Calculate ion concentrations in solutions of strong electrolytes. (Sect. 4.1) Prob: 29, 33, 35, 42 12. Differentiate between strong and weak acids, strong and weak bases, and salts. (Sect. 4.2) Prob: 46, 47 13. Solve stoichiometry problems for reactions between substances that are in solution. (Sect. 4.3) Prob: 113 14. Describe how a titrations is performed, and carry out calculations related to titrations. (Sect. 4.6) Prob: 76 b, 77, 83, 92, 102 Laboratories to be performed: 1. Determination of the concentration of an unknown solution using colorimetry. 2. Determination of the mass percent of acetic acid in vinegar (includes the standardization of the sodium hydroxide solution). 3. Oxidation-reduction titration: Determination of the mass percent of oxalate in an unknown sample. Unit 3: Understanding Chemical Reactions The student will 1. Have a concrete understanding of the solubility rules. 2. Apply the solubility rules when writing a balanced net ionic equation for a given chemical reaction. 3. Describe spectator ions in a solution, and be able to write net ionic equations for reactions that occur in solution. (Sect. 4.3) 4. Differentiate between the major types of chemical reactions, including but not limited to acid/base, precipitation, decomposition, synthesis, combustion, and oxidation/reduction. 5. Predict the products of the major types of chemical reactions. 6. Assign oxidation numbers to elements in compounds or ions. (Sect. 4.4) Prob: 68 assign oxidation #’s only. Do not balance. 7. Define and recognize oxidizing and reducing agents. (Sect. 4.4) Prob: 70 8. Balance redox reactions that occur in acidic, basic or neutral solutions. (Sect. 4.4 & 18.2) Prob in Ch 18: 27, 30 Laboratories to be performed: 1. The Copper Cycle: Various Chemical Reactions of Copper Unit 4: Chemical Kinetics Chapter 13: Chemical Kinetics: Rates and Mechanisms of Chemical Reactions The student will 1. Describe the kinetics of chemical reactions in terms of the collision theory and the transition state theory. (Sect. 13.7) Prob: 51 2. Construct reaction profiles and identify activation energies, transition states, and reaction enthalpies on these profiles. (Sect. 13.7) Prob: 53, 54 3. Predict the influence of different experimental factors on reaction rates. (Sect. 13.1) 4. Express reaction rates in terms of concentrations of reactants and/or products and use experimental data to determine the average rate and/or instantaneous rate of reaction. (Sect. 13.2 – 13.3) Prob: 23, 24, 25 a & b, 33, 75 5. Use the method of initial rates to determine the rate law and rate constant of a chemical reaction, as well as apply the rate law to determine the rate of reaction at different reactant concentrations. (Sect. 13.4) Prob: 32 6. Relate concentration of reactants and time using the integrated for zero-, first-, and second-order reactions. (Sect. 13.5 -13.6) Prob: 27, 30, 35, 40, 42, 43, 46 7. Explain the meaning of half-life, and derive and use equations for half-life of zero-, first-, and second-order reactions. (Sect. 13.5 -13.6) Prob: 39, 45, 50 8. Calculate the effects of temperature on rates of reaction, using the Arrhenius equation. (Sect. 13.8) Prob: 56, 57, 80 9. Describe the properties of elementary reactions and rate determining step in a reaction mechanism. (Sect. 13.9) 10. Deduce rate laws from simple reaction mechanisms, and establish the plausibility of reaction mechanisms based on kinetic data. (Sect. 13.9) Prob: 62, 64, 65, 68, 85 11. Describe the effects of a catalyst on the rate of reaction. (Sect. 13.10) Laboratories to be performed: 1. Determination of the Order of a Chemical Reaction Unit 5: Chemical Equilibrium: An introduction Chapter 14: Chemical Equilibrium The student will 1. Describe the nature of dynamic equilibrium in terms of forward and reverse processes. (Sect. 14.1) 2. Write the equilibrium constant expression for reversible chemical reactions. (Sect. 14.2) Prob: 10, 12 3. Derive equilibrium constants for other reactions related to a given reaction with a known equilibrium constant. (Sect. 14.3) Prob: 22, 24, 30 4. Derive the equilibrium constant for an overall reaction obtained by combining two or more reactions and their equilibrium constants. (Sect. 14.3) Prob: 26 5. Distinguish between Kc and Kp, and obtain the value for one of them from a known value of the other. (Sect. 14.3) Prob: 20, 23 6. Determine the value of the equilibrium constants from experimental data. (Sect. 14.5) Prob: 32, 34, 57, 58 7. Given initial concentrations, calculate equilibrium concentrations or partial pressure using the given equilibrium constant. (Sect. 14.5) Prob: 60, 62, 69, 72, 81 8. Describe nonequilibrium conditions through a reaction quotient, and use this value to predict the direction in which a net reaction must occur to achieve equilibrium. (Sect. 14.3) Prob: 49, 52, 66 9. Use LeChâtelier’s principle to predict shifts in equilibria following changes in concentration, temperature, pressure, or the addition of a catalyst. (Sect. 14.4) Prob: 39, 43, 75 10. Use the van’t Hoft equation to determine the value of the equilibrium constant at a different temperature. (Sect. 17.7 (note different chapter) Prob: 88 Chapter 16: More Equilibria in Aqueous Solutions: Slightly Soluble Salts and Complex ions The student will 11. Write expressions for Ksp describing solubility of slightly soluble salts. (Sect. 16.1) Prob: 19, 20, 21, 22 12. Calculate Ksp values from the molar solubility of salts and vice versa. (Sect. 16.2) Prob: 25, 28, 30, 32, 35 13. Predict and calculate the effect of addition of a common ion on the solubility of a salt. (Sect. 16.3) Prob: 38, 40, 41, 43 14. Use the reaction quotient to determine whether a precipitation will occur when two ions are mixed . (Sect. 16.4) Prob: 47, 49, 50, 53, 58, 60 15. Predict and calculate the effect of pH on the solubilities of salts. (Sect. 16.5) Prob: 64, 66, 90, 94, 100 16. Using formation constants of complex ions, determine the concentrations of ions in solutions containing complex ions. (Sect. 16.6) Prob: 68, 70, 71, 73, 76 Laboratories to be performed: 1. Examination of LeChâtelier’s Principle in a Microplate System. 2. Determining the Ksp of Calcium Hydroxide. Unit 6: Acid/Base Equilibria and Solubility Equilibria Chapter 15: Acids, Bases, and Acid/Base Equilibria The student will 1. Define acids and bases according to the Arrhenius, Brønsted-Lowry, and Lewis theories. (Sect. 15.1 & 15.11) 2. Write equilibrium constant expressions for Brønsted-Lowry acid and bases, and describe what is meant by conjugate acid/base pairs. (Sect. 15.1) Prob: 24, 26, 28 3. Predict the influences of molecular structure on the relative strengths of acids and bases. (Sect. 15.2) Prob: 31, 32 4. Relate equilibrium values of [H3O+] and [OH-] to another by using the ion product of water. (Sect. 15.3) Prob: 38 5. Determine values of conjugate acid and conjugate base ionization constants through their relationship to Kw. (Sect. 15.3) Prob: 67 a & b 6. Calculate the pH value for strong acids, strong bases, weak acids, and weak bases. (Sect. 15.3 – 15.5) Prob: 42, 47, 48, 49, 52, 54, 56, 59 7. Calculate the pH of salt solutions (Sect 15.6) Prob: 63, 65, 67 c; 6, 77 8. Predict and calculate the effect of pH on the solubility of salts. (Sect. 16.5) Prob:( In Chapter 16) 48, 64, 66, 90, 94, 100 9. Describe the common ion effect, and predict the influence on acid/base equilibria of adding common ions to a solution. (Sect. 15.7) Prob: 71 10. Calculate equilibrium concentrations and pH values for solutions of weak acids and weak bases containing common ions. (Sect. 15.8) Prob: 73, 74, 75, 76 11. Describe the action of a buffer solution, and calculate equilibrium concentrations and pH values of such solutions using either Ka or Kb values or the Henderson-Hasselbalch equation. (Sect. 15.8) Prob: 81, 82, 83, 84, 103, 122 12. Calculate the effect on the pH produced by adding a strong acid or strong base to a buffer solution, and describe the pH range and buffer capacity of the solution. (Sect. 15.8) Prob: 85, 86, 88, 106 13. Describe the function of an acid/base indicator and the relationship between pH and the color of the indicator, select appropriate indicators for a given pH range, and determine the color of an indicator in a given solution. (Sect. 15.9) Prob: 90 14. Describe the regions of a titration curve in terms of equilibria that exist at different points along curve. Calculate the pH values a various points along the curve, and determine the Ka value of the acid using the curve. (Sect. 15.10) Prob: 97, 99, 128 Laboratories to be performed: 1. Determination of the Acid Ionization Constant of a Weak Acid. 2. Preparation of a Buffer Solution at a Given pH. 3. Determination of the Equilibrium Constant of an Indicator. Unit 7: Thermodynamics Chapter 6: Thermochemistry The student will 1. Define and distinguish between kinetic energy, potential energy, and internal energy. (Sect. 6.1 – 6.2) 2. Identify systems and surroundings and describe the three types of thermodynamic systems. (Sect. 6.2) Prob: 1 3. Describe the transfer of heat between a system and its surroundings and what is meant by thermal equilibrium. (Sect. 6.4) 4. State and apply the law of conservation of energy and the first law of thermodynamics. (Sect. 6.3) 5. Describe what is meant by a state function. (Sect. 6.4) Prob: 4 6. Define and apply the terms exothermic and endothermic reaction. (Sect. 6.4) 7. Perform calculations involving enthalpy changes in chemical reactions. (Sect. 6.4) Prob: 28, 30, 34, 35, 38, 40, 42, 92 (use info found in 79) 8. Perform calorimetric calculations. (Sect. 6.5) Prob: 46 a, 49, 51, 54, 56, 57, 60, 99 9. State Hess’s law of constant summation, explain its significance, and apply it. (Sect. 6.6) Prob: 67, 70 10. Understand the meanings of standard state, the standard enthalpy of reaction, and the standard enthalpy of formation. (Sect. 6.7) 11. Use bond energies to calculate the enthalpy change for a chemical reaction (Sect. 9.10) Prob: (in Ch. 9) 69, 70 12. Use standard enthalpy of formations to determine standard enthalpy changes for chemical reactions. (Sect. 6.7) Prob: 72 b & c, 74, 97 Chapter 17: Thermodynamics: Spontaneity, Entropy, and Free Energy The student will 13. Describe the differences between spontaneous and nonspontaneous processes. (Sect. 17.2) 14. Define entropy and describes its relationship with spontaneous processes. (Sect. 17.3) Prob: 24 15. State the second and third law of thermodynamics. (Sect. 17.3) 16. Predict the relative entropy change for a given process. (Sect. 17.3) Prob: 19, 26 17. Relate the value of entropy and enthalpy to the spontaneity of a given process. (Sect. 17.4) 18. Determine Gibbs free energy and relate the value to the spontaneity of a given process. (Sect. 17.4 – 17.5) Prob: 30, 36, 38, 40 19. Calculate the equilibrium constant from Gibbs free energy. Relate the magnitude and sign of Gibbs free energy to the position of equilibrium of a given process. (Sect. 17.6) Prob: 45, 49, 54, 55, 57 20. Calculate Gibbs free energy for a given process under nonstandard conditions. (Sect. 17.6) Prob: 44, 61, 70, 82 Laboratories to be performed: 1. Determining the Enthalpy change of a Chemical Reaction using Hess’s Law. Unit 8: Electrochemistry Chapter 18: Electrochemistry The student will 1. Describe the basis of electrochemical cells in terms of oxidation-reduction reactions. (Sect. 18.1 – 18.2) Prob: 28, 30 2. Describe a voltaic cell and the parts that make up a voltaic cell. (Sect. 18.3) 3. Use standard reduction potentials to calculate standard cell potential for a voltaic cell. (Sect. 18.4) Prob: 33, 36, 38, 40, 42 4. Relate the standard cell potential to the spontaneity of the voltaic cell. (Sect. 18.5) Prob: 44, 46, 48 5. Calculate Gibbs free energy and the equilibrium constant for voltaic cells at standard conditions. (Sect. 18.5); Prob: 52, 54, 56, 107 6. Determine cell potentials under nonstandard conditions using the Nernst equation. (Sect. 18.6) Prob: 58, 60, 61, 64, 97 7. Use the Nernst equation to calculate potentials of concentration cells. (Sect. 18.6) Prob: 102 8. Describe the relationship between voltaic cells and various kinds of batteries. (Sect. 18.7) 9. Describe electrolysis and calculate the amount of substances consumed or produced during electrolysis reactions. (Sect. 18.9 – 18.10) Prob: 77, 80, 82, 83, 91 Laboratories to be performed: 1. Using Concentration Cells to Determine the Concentration of an Unknown Solution. Unit 9: Nuclear Chemistry Chapter 19: Nuclear Chemistry The student will 1. Describe the basic nature of α, β−, and γ radiation. (Sect. 19.1) 2. Write and balance nuclear equations involving the five different decay modes of radioactive nuclides. (Sect. 19.1) ` Prob: 22, 23 a, 24 a & b, 26 3. Predict the products of a radioactive decay series. (Sect. 19.2) Prob: 23 b & c, 24 c 4. Calculate decay rates and half-lives of radioactive isotopes. (Sect. 19.3) Prob: 29, 30, 31, 38, 39, 84, 88 a, 90 a & b 5. Predict the nuclear stability of specific isotopes, using magic numbers and the belt of stability. (Sect. 19.6) Prob: 59, 60 6. Relate mass changes to energy changes in nuclear equations through calculations involving E = mc2. (Sect. 19.7); Prob: 49, 52, 54, 55 7. Describe the use of nuclear fission and nuclear fusion in energy production. (Sect. 19.8) Prob: 66 Unit 10: Electronic Structure Chapter 7: Atomic Structure The student will 1. Describe how a cathode ray tube works and the major results of Thomson’s experiments. (Sect. 7.1) Prob: 24 2. Describe how Millikan’s experiments were used to determine the charge of the electron. (Sect. 7.1) Prob: 70 3. Explain the results of Rutherford’s gold foil experiment and how these results led to Rutherford’s nuclear model of the atom. (Sect. 7.2) 4. Explain the differences between Thomson’s plum pudding model of the atom and Rutherford’s nuclear model of the atom. (Sect. 7.2) 5. Describe the properties of a proton and neutron and how they were discovered. (Sect. 7.3) 6. Calculate the wavelength, frequency, and/or energy of electromagnetic radiation. (Sect. 7.5-7.6) Prob: 38, 39, 43, 74 7. Describe the difference between continuous and line spectra. (Sect. 7.5) 8. Describe the Bohr model of the hydrogen atom. (Sect. 7.7) Prob: 77, 78 9. Calculate the energy of an electron in a given energy level of a hydrogen atom and the energy differences in the levels. (Sect. 7.7) Prob: 48 10. Describe the wave properties of matter and the nature of wave mechanics and wave function. (Sect. 7.8) Prob: 54, 56 11. Have an understanding of quantum numbers, their possible values, and their relationship to atomic orbitals. (Sect. 7.9) Prob: 61, 62, 65, 66 12. Describe the meaning of atomic orbitals and the shapes of s, p, and d orbitals. (Sect. 7.9) Chapter 8: Electron Configurations, Atomic Properties, and the Periodic Table The student will 13. Explain the origin of energy differences between the orbitals of a hydrogen atom and those of a multielectron atom and between different subshells within a principle shell of a multielectron atom. (Sect. 8.1) 14. Write electron configurations using spdf notation and orbital diagrams of atoms and ions. (Sect. 8.2) Prob: 28 b, e, g, & i, 30 b, c, & d, 32 b & d, 35 b & c, 38 15. Apply the Pauli exclusion principle, Hund’s rule, and the Aufbau, principle to the writing of electron configurations. (Sect. 8.3-8.4) Prob: 22, 24, 26 16. Deduce electron configurations of atoms based on their position in the periodic table. (Sect. 8.5) Prob: 44, 75, 76 Unit 11: Periodic Properties, Chemical Bonding, and Molecular Geometries Chapter 8: Electron Configurations, Atomic Properties, and the Periodic Table The student will 1. Predict the magnetic properties of atoms and ions. (Sect. 8.6) Prob: 46 2. Define atomic and ionic radius, ionization energy, electron affinity, and electronegativity, and describe the trends of each as they relate to the periodic table. (Sect. 8.7) Prob: 52, 54, 56, 58, 59, 62 3. Describe the characteristic properties of metals, metalloids, and nonmetals. (Sect. 8.8) Prob: 68 4. Explain trends in macroscopic properties of elements. (Sect. 8.9) Prob: 70 Chapter 9: Chemical Bonds The student will 5. Describe the energetics of chemical bond formation. (Sect. 9.5) Prob: 27, 30 6. Cite the characteristics of ionic and covalent bonds. (Sect. 9.2 - 9.3) Prob: 23 7. Describe the nature of multiple bonds in terms of shared electron pairs. (Sect. 9.6) 8. Describe the influence of electronegativity on chemical bonding and bond polarity. (Sect. 9.7) Prob: 40, 42 9. Draw Lewis dot structures for molecules and ions, including any exceptions. (Sect. 9.8 – 9.9) Prob: 36, 38, 56, 62, 101 10. Calculate formal charges and use them in assessing the plausibility of Lewis dot structures. (Sect. 9.8) Prob: 46, 48 11. Draw resonance structures when appropriate and describe the resonance hybrid. (Sect. 9.8) Prob: 50 12. Define bond length, describe the general relationship between bond length and bond order, and predict bond lengths for the Lewis structures. (Sect. 9.10) Prob: 64, 67, 91 Chapter 10: Bonding Theory and Molecular Geometry The student will 13. Identify the geometrical shapes of molecules and polyatomic ions according to the principles of the VSEPR theory. (Sect. 10.1) Prob: 25, 26, 27, 28, 31 14. Predict the bond angles within a molecule and influence of lone pairs on bond angles. (Sect. 10.1) Prob: 33, 35 15. Describe the shape of molecules having more than one central atom. (Sect. 10.1) Prob: 36 16. Define a dipole moment and categorize molecules as polar or nonpolar. (Sect. 10.2) Prob: 38, 40, 77 17. Describe the valence bond theory of covalent bond formation. (Sect. 10.3) 18. Describe hybridization of atomic orbitals in terms of orbital energy and spatial orientation, and predict the central atom hybridization and atomic orbital overlap associated with covalent bond formation. (Sect. 10.4) Prob: 45, 48, 50, 51, 84 19. Describe the hybridization schemes that lead to multiple bond formation, and identify the sigma and pi bonds present in a given molecular structure. (Sect. 10.5) Prob: 56 20. Describe cis-trans isomers and the bonding that lead to them. (Sect. 10.5) Prob: 57 21. Draw and name simple aromatic structures. (Sect. 10.9) Prob: 69, 70, 71, 72 Unit 12: Intermolecular Forces, Liquids, and Solids Chapter 11: States of Matter and Intermolecular Forces The student will 1. Describe the differences between dispersion, dipole-dipole, dipole-induced-dipole forces, and hydrogen bonding. (Sect. 11.5 – 11.6) 2. Determine the type of intermolecular force present in a given compound. (Sect. 11.5 – 11.6) 3. Describe the origin of intermolecular forces, and predict physical properties of the substances from the nature of their intermolecular forces. (Sect. 11.5 - 11.6) Prob: 49, 51, 55, 58 4. Describe the distinguishing characteristics of the three states of matter, and identify and describe the six main types of phase changes. (Sect. 11.1) 5. Perform calorimetric calculations related to phase changes. (Sect. 11.2 - 11.3) Prob: 38, 39 6. Define vapor pressure, boiling point, melting point, surface tension, and viscosity, and relate each to the type of force found in a substance. (Sect. 11.2, 11.3, 11.7) Prob: 60, 74, 92 7. Perform calculations related to vapor pressure. (Sect. 11.2) Prob: 29, 31, 73, 83 8. Use the Clausius-Clapeyron equation to determine vapor pressures of a substance at different temperatures. (Sect. 17.7) Prob: (in Ch. 17) 63, 64 9. Construct and interpret phase diagrams. (Sect. 11.4) Prob: 44, 45, 93 10. Compare the properties of ionic, molecular, metallic, and network-covalent solids. (Sect. 11.8 – 11.9) 11. Discuss the organization of atoms/molecules in a crystalline solid in terms of unit cells and close-packed crystal structure. (Sect. 11.10) Prob: 64 Chapter 12: The Physical Properties of Solutions The student will 12. Explain the effects of a solute particle on the freezing point, boiling point, and vapor pressure of a pure liquid. (Sect. 12.6 – 12.7) 13. Discuss the differences between electrolytic and nonelectrolytic solutes using the van’t hoft factor. (Sect. 12.9) 14. Calculate the vapor pressure of a solution using Raoult’s law, and explain variations to Raoult’s law found in non-ideal solutions. (Sect. 12.6) Prob: 55, 56 15. Describe the process of fractional distillation in relation to the vapor pressures of solution components. (Sect. 12.6) 16. Calculate changes in freezing point and boiling point of a liquid due to the presence of a solute. (Sect. 12.7); Prob: 57, 60, 75, 87, 91 17. Calculate the osmotic pressure of a solution. (Sect. 12.8) Prob: 69, 70 18. Calculate the molar mass of an unknown substance from freezing point, boiling point, or osmotic pressure data. (Sect. 12.7 – 12.8) Prob: 62, 63, 97 19. Discuss ion pairing and its effect on a given colligative property. (Sect. 12.9) Prob: 71, 72 Laboratories to be performed: 1. Determining the Molar Mass of an Unknown Substance by Freezing Point Depression