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.

Write the formulas for the following: Name each of the following:

1. Barium Sulfate 1. CuSO4

2. Ammonium Chloride 2. PCl3

3. Chlorine Monoxide 3. Li3N

4. Silicon Tetrachloride 4. BaSO4

5. Magnesium Fluoride 5. N2F4

6. Sodium Oxide 6. KClO4

7. Sodium Peroxide 7. NaH

8. Copper (I) Oxide 8. (NH4)Cr2O7

9. Zinc Sulfide 9. Ba(NO2)2

10. Potassium Carbonate 10. Sr3P2

11. Manganese (II) Phosphate 11. Mg(OH)2

12. Silicon Tetrabromide 12. Al2S3

13. Lead (II) Acetate 13. AgBr

8 Write the formulas for the following: Name each of the following: 1. Sodium Permanganate 1. P4O10

2. Lithium Oxalate 2. Pb(C2H3O2)4

3. Potassium Cyanide 3. CaI2

4. Iron (III) Hydroxide 4. MnO2

5. Silicon Dioxide 5. Li2O

6. Nitrogen Trifluoride 6. FeI3

7. Chromium (III) Oxide 7. Cu3PO4

8. Calcium Chlorate 8. PCl5

9. Sodium Thiocyanate 9. NaCN

10. Aluminum Nitride 10. IF3

9 Chemical Bonding – Rules for Naming Acids Acids are covalent compounds that ionize or form ions in water:

+ - HCl + H2O → H3O + Cl (chloride) Acid names are linked to the anion formed when the acid reacts with water.

1. When the name of the anion (negative ion) ends in -ide, the acid name begins with the prefix hydro- and the stem of the anion has the suffix -ic. The renames anion is followed by the word acid.

Pattern: Hydro ic acid Examples: HCl = Hydrochloric acid [Cl- = chloride] HCN = hydrocyanic acid [CN- = cyanide]

2. When the name of the anion ends in -ite, the acid name is the stem of the anion with the suffix -ous. This renamed anion is followed by the word acid.

Pattern: ous acid - Examples: HNO2 = Nitrous acid [NO2 = nitrite] 2- H2SO3 = Sulfurous acid [SO3 = sulfite]

3. When the name of the anion ends in -ate, the acid name is the stem of the anion with the suffix -ic. This renamed anion is followed by the work acid.

Pattern: ic acid - Examples: HNO3 = nitric acid [NO3 = nitrate] 2- H2SO4 = Sulfuric acid [SO4 = sulfate]

10 Acids Names & Formulas Exercise: Complete the page using the rules on page 10.

Name the following acids:

1. HBr 4. HC2H3O2 7. HClO

2. H3PO4 5. HI 8. H2C2O4

3. HClO4 6. HIO3

Give the formula for the following acids: 1. Hydrofluoric acid 5. Chlorous acid 2. Bromic acid 6. Phosphorous acid 3. Chloric acid 7. Selenic acid 4. Carbonic acid

11 Reactions – Solubility Rules (for aqueous solutions)

Ion Solubility % ��$ % �#�$�# Always Soluble % ���! % ���$ Alkali Metals Always Soluble " ��! ��% % " #" #" �� Soluble except with �� , �� , ��# �% Soluble except with ��#", ��#", #% # ��! ��#",��#", ��#" ��% Insoluble except with ��#", ��#", ��#" % ��! �#% Insoluble except with Alkali metals and #% " ��$ ��! #% ��$

Soluble acids = HCl, HBr, HI, ���$,

Acids ����!, �#��! Insoluble acids = all other acids

12 Questions: 1. What is meant by the term solution?

2. What does it mean when we say a compound is soluble or insoluble?

3. What does the term “aqueous” mean?

4. Are the solubility rules different for compounds mixed with alcohol or other liquids? 5. What elements are included in the alkali metal group mentioned in the solubility rules?

Assignment: It is important to know the rules for solubility of ionic compound (and acids) during many of the units in AP Chemistry. Make flashcards of the solubility rules: Write the formula for the ion on one side of the flashcard and the rule for the solubility on the other side of the card.

13 Applying Solubility Rules

Exercise: Use the examples below to complete the chart on the following page.

Compound Soluble or Insoluble? Ions present in solution? KCl - Potassium Chloride Soluble K+ and Cl- Solubility Rule

+ - + 2+ 2+ Alkali metal K is always soluble or Cl (chloride ion) is soluble except with Ag , Pb , and Hg2

Compound Soluble or Insoluble? Ions present in solution?

2+ - Cu(NO3)2 – Copper (II) nitrate Soluble Cu and NO3 Solubility Rule Nitrate is always soluble

Compound Soluble or Insoluble? Ions present in solution? AgCl – Silver Chloride Insoluble None Solubility Rule

+ 2+ 2+ chloride ion is soluble except with Ag , Pb , and Hg2

Compound Soluble or Insoluble? Ions present in solution?

Zn(OH)2 Insoluble None Solubility Rule Hydroxide ion is insoluble except with Ca2+, Sr2+, and Ba2+

14 Applying Solubility Rules

Exercise: For each compound below determine whether the compound is soluble or insoluble in water. Indicate which ions, if any, would be present in an aqueous solution.

Compound Soluble or Insoluble? Ions present in solution? NaCN

NH4NO3

CuSO4 Sulfurous Acid

FeI3

KC2H3O2

HNO3 Copper (I) Chlorate

Al2S3 Lead (II) Bromide

(NH4)2Cr2O7

Na3N Zinc Sulfide

BaSO3 Potassium Chlorate NaH

Mg(OH)2 Calcium Iodide

Cu3(PO4)2

K2CO3

ZnSiO3

15 Reactions – Chemical Equations

Diatomic Molecules: H2, O2, N2, F2, Cl2, Br2, and I2 (memorize these) Examples of five types of reactions (equations are unbalanced) 1. Composition/Synthesis/Combination a. Example: Lithium burns in Oxygen gas i. Element + element → compound (check to see if the elements involved are diatomic)

ii. Li + O2 → Li2O (use charges of the ions to determine the formula of the compound) 2. Decomposition a. Example: calcium oxide is heated i. Compound → element + element (could be element and compound or compound and another compound) (heat applied to the compound triggers the decomposition of the compound)

ii. CaO → Ca + O2 3. Single Replacement/Displacement a. Example: Lithium reacts with sodium nitrate i. Element + Compound → compound + element the three ions in the reactants switch

ii. Li + NaNO3 → LiNO3 + Na (partners to form a new compound (LiNO3) and different free element (Na)) 4. Double Replacement/Displacement a. Example: Calcium oxide reacts with barium sulfate i. Compound + Compound → Compound + Compound

ii. CaO + BaSO4 → CaSO4 + BaO (the four ions in the reactant compounds switch partners as to form two new products) 5. Combustion a. Example: Hydrocarbon react with oxygen i. Hydrocarbon + Oxygen → Carbon dioxide + water

ii. CH4 + O2 → CO2 + H2O (a hydrocarbon contains C, H and may contain O also)

16 Chemical Equations

Exercise: For each reaction below: • Identify the type of reaction • Predict and write the products • Balance the reaction Hint: When writing these reactions, ignore all of the information about heat, bubbling, and mixing details. These are excess words used to make complete sentences. Simply locate the chemical formulas or names in the description to find the reactants for the reaction. Example: solutions of silver nitrate and magnesium iodide are combined.

Double replacement – 2 AgNO3 + MgI → 2 AgI + Mg(NO3)2

1. Ammonium sulfate reacts with barium nitrate.

2. Zinc metal is added to a solution of copper (II) chloride.

3. Propane gas (C3H8) is burned in excess oxygen.

4. Magnesium and nitrogen gas are heated together.

5. Chlorine gas is bubbled thorough a solution of sodium bromide.

6. Solutions of Lead (II) nitrate and calcium iodide are mixed.

7. Sulfuric acid is combined with sodium hydroxide.

8. Isopropyl alcohol (C3H7OH) is burned in oxygen.

9. Iron metal shavings are added to hydrochloric acid.

10. Ammonium hydroxide solution is added to potassium sulfite.

Be sure all equations are balanced.

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