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AP Chemistry Summer Assignment AP Chemistry Summer Assignment If you have chosen to take Advance Placement Chemistry, you should have a very good background in chemistry from Honors Chemistry I or Pre-IB Chemistry. Advance Placement Chemistry is a college level course covering topics including electrochemistry, equilibrium, chemical kinetics and thermodynamics. Rather than memorizing how to do a particular type of problem, you must really understand the chemistry and be able to apply it to different situations. Because of the amount of material, we must cover before the AP exams in May students must complete much of the work outside of class. Homework will include practice problems, sample AP questions and reading assignments from the textbook. But, with hard work, you will not only be successful in the AP Chemistry exam and course, but also be prepared for college level course work. Like most AP classes, AP chemistry comes with a summer assignment. Previous AP students have helped design this assignment – it is what they think is important to review and know before starting class in the fall. The assignment will be collected the second Friday of the school year. Do not procrastinate during the summer! You will need time to complete the different parts of this assignment, memorize some items and review before school starts. So, schedule your work during the summer! Resources: • Review book – there are many out there. The following are just suggestions. o “AP Chemistry” 8th edition or later – Barron’s o “AP Chemistry” Crash Course Materials for Class: • Composition Notebook for lab work – Graphing (Might be provided by the school) • Scientific Calculator (I have a class set, but highly suggested for doing work at home.) • School Issued laptop • Composition notebook for notes and practice problems 1 AP Chemistry Course Content 1. Atomic Structure and Properties a. Moles and Molar Mass b. Mass Spectroscopy of Elements c. Elemental Composition of Pure Substances d. Composition of Mixtures e. Atomic Structure and Electron Configuration f. Photoelectron Spectroscopy g. Periodic Trends h. Valence Electrons and Ionic Compounds 2. Molecular and Ionic Compound Structure and Properties a. Types of Chemical Bonds b. Intramolecular Force and Potential Energy c. Structure of Ionic Solids d. Structure of Metals and Alloys e. Lewis Diagrams f. Resonance and Formal Charge g. VSEPR and Bond Hybridization 3. Intermolecular Forces and Properties a. Intermolecular Forces b. Properties of Solids c. Solids, Liquids, and Gases d. Ideal Gas Law e. Kinetic Molecular Theory f. Deviation from Ideal Gas Law g. Solutions and Mixtures h. Representations of Solutions i. Separation of Solutions and Mixtures Chromatography j. Solubility k. Photoelectric Effect l. Beer-Lambert Law 4. Chemical Reactions a. Introduction to Reactions b. Net Ionic Equations c. Representations of Reactions d. Physical and Chemical Changes e. Stoichiometry f. Introduction to Titration g. Types of Chemical Reactions h. Introduction to Acid-Base Reactions i. Oxidation-Reduction (Redox) Reactions 5. Kinetics a. Reaction Rates b. Introduction to Rate Law c. Concentration Changes Over Time d. Elementary Reactions e. Collision Model f. Reaction Energy Profile g. Introduction to Reaction Mechanisms h. Reaction Mechanism and Rate Law i. Steady-State Approximation j. Multistep Reaction Energy Profile k. Catalysis 2 6. Thermodynamics a. Endothermic and Exothermic Processes b. Energy Diagrams c. Heat Transfer and Thermal Equilibrium d. Heat Capacity and Calorimetry e. Energy of Phase Changes f. Introduction of Enthalpy of Reaction g. Bond Enthalpies h. Enthalpy of Formation i. Hess’s Law 7. Equilibrium a. Introduction to Equilibrium b. Direction of Reversible Reactions c. Reaction Quotient and Equilibrium Constant d. Calculating the Equilibrium Constant e. Magnitude of Equilibrium Constant f. Properties of the Equilibrium Constant g. Calculating Equilibrium Concentrations h. Representations of Equilibrium i. Introduction to Le Châtelier’s Principle j. Reaction Quotient and Le Châtelier’s Principle k. Introduction to Solubility and Equilibria l. Common-Ion Effect m. pH and Solubility n. Free Energy of Dissolution 8. Acids and Bases a. Introduction to Acids and Bases b. pH and pOH of Strong Acids and Bases c. Weak Acids and Base Equilibria d. Acid-Base Reactions and Buffers e. Acid-Base Titrations f. Molecular Structure of Acids and Bases g. pH and pKa h. Properties of Buffers i. Henderson-Hasselbach Equation j. Buffer Capacity 9. Applications of Thermodynamics a. Introduction to Entropy b. Absolute Entropy and Entropy Change c. Gibbs Free Entropy and Thermodynamic Favorability d. Thermodynamic and Kinetic Control e. Free Energy and Equilibrium f. Coupled Reactions g. Galvanic (Voltaic) and Electrolytic Cells h. Cell Potential and Free Energy i. Cell Potential Under Nonstandard Conditions j. Electrolysis and Faraday’s Law 3 Summer Assignment Check off each part of the assignment as you complete it and be sure to bring the appropriate items for collection on the second Friday of the school year. Do not just complete the material in the summer assignment – make sure that you know and understand the material. • Chemical Bonding – Cations & Anions o Pages 5 – 6 of this Packet § Complete the exercise and question sections. § Play the “Plainfield AP Chemistry Polyatomic Review” at https://www.youtube.com/watch?v=w3zJIKHsjZI § Prepare flashcards as described in the assignment and study the ion names and formulas. • Chemical Bonding – Compound Formulas & Names o Page 7 of this packet: § Go over examples of ionic and covalent compound nomenclature. o Pages 8-9 - complete the Compound Nomenclature assignment • Chemical Bonding – Rules for Naming Acids o Page 10 of this packet § Go over the examples of acid naming rules. o Page 11 - complete the exercise. • Reactions – Solubility Rules o Page 12 - Answer the questions o Page 13 – Prepare the flashcards as described in the assignment and study the solubility rules. o Page 14 – Use the examples provided to complete the examples provided. o Page 15 – Compete the exercise. • Reactions – Chemical Equations o Page 16 – Go over the types of equations in the examples provided. o Page 17 - Complete the exercise. Be sure to balance each equation. Bring these items with you to class the second Friday of school: 1. Flashcards for Common Ions. 2. Flashcards for Solubility Rules 3. Completed summer Packet – Staple or paperclip the entire packet together to turn in. 4 Chemical Bonding – Cations & Anions Common Ions with One Oxidation State Name Formula Name Formula Lithium Li+1 Magnesium Mg+2 Sodium Na+1 Calcium Ca+2 Potassium K+1 Strontium Sr+2 Silver Ag+1 Barium Ba+2 Aluminum Al+3 Zinc Zn+2 Hydrogen H+1 or H-1 Nitride N-3 Oxide O-2 Sulfide S-2 Fluoride F-1 Chloride Cl-1 Bromide Br-1 Iodide I-1 Exercise: Locate the elements listed above on your periodic table. Color all atoms with a +1- state purple, atoms with a +2-state blue, atoms with a +3 state green, atoms with a -3 state yellow, atoms with a -2 state orange, and atoms with a -1 state red. Label the top of each column involved with the appropriate ion charge. Common Ions with more than One Oxidation State Name Formula Name Formula Copper (I) Cu+1 Mercury (I) Hg+1 Copper (II) Cu+2 Mercury (II) Hg+2 Iron (II) Fe+2 Chromium (I) Cr+1 Iron (III) Fe+3 Chromium (III) Cr+3 Manganese (II) Mn+2 Cobalt (II) Co+2 Manganese (III) Mn+3 Cobalt (III) Co+3 Tin (II) Sn+2 Lead (II) Pb+2 Tin (IV) Sn+4 Lead (IV) Pb+4 Questions: 1. Color each element listed above on your periodic table brown. 2. Is there a trend based on the oxidation state for each of the transition elements? Explain. 3. What does the Roman numeral indicate about the ion? 4. Why is it not necessary to give a Roman numeral when writing the name for silver and zinc ions? 5. Write the charges above the box on your periodic table for each of the transition elements listed above. 5 Polyatomic Ions A charged particle which has two or more atoms held together by covalent bonds. Name Formula Name Formula " % Ammonium ��! Acetate �#�$�# #% #% Oxalate �#�! Peroxide �# % % Hydroxide �� Permanganate ���! #% % Selenate ���! Bromate ���$ % #% Iodate ��$ Silicate ���$ % % Nitrate ��$ Cyanide �� % % Nitrite ��# Thiocyanate ��� $% #% Phosphate ��! Chromate ���! $% #% Phosphite ��$ Dichromate ��#�& #% % Carbonate ��$ Perchlorate ���! % % Bicarbonate ���$ Chlorate ���$ #% % Sulfate ��! Chlorite ���# #% % Sulfite ��$ Hypochlorite ��� % Bisulfate ���! % Bisulfite ���$ Questions: 1. What do the prefixes hypo-, per-, bi-, and thio- mean in terms of polyatomic ions? 2. What is the significance of the suffixes -ate and -ite in terms of the formula of the ions? Assignment: Students taking the AP Chemistry exam are given several reference sheets to use during the test. Monoatomic and polyatomic ions are not included on those reference sheets, so students are required to memorize them. Make Flashcards of the ions from the list above. Write the ion formula with chare on one side of the flashcard and the name of the ion on the other side of the card. You can also use the free website www.quizlet.com to make, learn, and print your flashcards. 6 Chemical Bonding – Compound Formulas & Names Use the examples below to complete the questions on the following pages Ionic Compounds Metal + non-metal Or containing a polyatomic ion Total ions charges must equal 0 Potassium Oxide Calcium Nitrate Iron (III) Hydroxide 1+ 2- 2+ - 3+ 1- Ions = K O ions = Ca NO3 ions = Fe OH Formula = K2O Formula = Ca(NO3)2 Formula = Fe(OH)3 BaBr2 (NH4)2CO3 Cu2O Barium Bromide Ammonium Carbonate Copper (I) Oxide Covalent Compounds Non-metal + Non-metal Or acid (HX) Prefixes in name indicate number of each element Sulfur Trioxide Phosphorous Pentachloride Dinitrogen Tetraoxide SO3 PCl5 N2O4 BCl3 SCl6 OCl2 Boron Trichloride Sulfur Hexachloride Oxygen Dichloride 7 Compound Nomenclature Exercise: Complete this section by giving the correct formula or name for each compound.
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