Download SUMMER WORK AP Chemistry

SUMMER WORK
AP Chemistry
Summer 2015
Dr. Kathleen Gorski, Chair, Science Department
[email protected]
Topics equivalent to those in Sections 1-5 (see below) are expected to be mastered prior to the start of the school year. Pay
special attention to the solubility rules, and be sure to know the common monatomic and polyatomic ions. Topics in
sections 6 – 8, if covered, are typically in less depth during a first year high school chemistry class, but should be familiar
to students. Less than two weeks are spent reviewing this material at the beginning of the school year. This material is
covered in any secondary school chemistry class, so students should already be familiar with it. The summer assignment is
meant to reinforce these topics as well as provide a brief introduction of simple related topics.
All studies should be performed in English. All work done in class will be in English and students who do not prepare
correctly will be at a great disadvantage.
The summer assignment will be collected at the beginning of the first class of the year. The assignment is a requirement for
enrollment in AP Chemistry.
1. Introduction: Matter and Measurement
• CLASSIFICATIONS OF MATTER: States of Matter, Pure Substances, Elements, Compounds, Mixtures
• PROPERTIES OF MATTER: Physical and Chemical Changes, Separation of Mixtures
• UNITS OF MEASUREMENT: SI Units, Length and Mass, Temperature, Derived SI Units, Volume, Density
• UNCERTAINTY IN MEASUREMENT: Precision and Accuracy, Significant Figures, Significant Figures in
Calculations
• DIMENSIONAL ANALYSIS: Using Two or More Conversion Factors, Conversions Involving, Volume
2. Atoms, Molecules, and Ions
• THE ATOMIC THEORY OF MATTER
• THE DISCOVERY OF ATOMIC STRUCTURE: Cathode Rays and Electrons, Radioactivity, The
Nuclear,Model of the Atom
• THE MODERN VIEW OF ATOMIC STRUCTURE: Atomic Numbers, Mass Numbers, and Isotopes
• ATOMIC WEIGHTS: The Atomic Mass Scale, Atomic Weight
• THE PERIODIC TABLE
• MOLECULES AND MOLECULAR COMPOUNDS: Molecules and Chemical Formulas, Molecular and
Empirical Formulas, Picturing Molecules
• IONS AND IONIC COMPOUNDS: Predicting Ionic Charges, Ionic Compounds
• NAMING INORGANIC COMPOUNDS: Names and Formulas of Ionic Compounds, Names and Formulas of
Acids , Names and Formulas of Binary Molecular Compounds
• SOME SIMPLE ORGANIC COMPOUNDS: Alkanes, Some Derivatives of Alkanes
3. Stoichiometry: Calculations with Chemical Formulas and Equations
• CHEMICAL EQUATIONS: Balancing Equations, Indicating the States of Reactants and Products
• SOME SIMPLE PATTERNS OF CHEMICAL REACTIVITY: Combination and Decomposition Reactions,
Combustion Reactions
• FORMULA WEIGHTS: Formula and Molecular Weights, Percentage Composition from Chemical Formulas
• AVOGADRO’S NUMBER AND THE MOLE: Molar Mass, Interconverting Masses and Moles, Interconverting
Masses and Numbers of Particles
• EMPIRICAL FORMULAS FROM ANALYSES: Molecular Formulas from Empirical Formulas, Combustion
Analysis
• QUANTITATIVE INFORMATION FROM BALANCED EQUATIONS
• LIMITING REACTANTS: Theoretical Yields
4. Reactions in Aqueous Solution
• GENERAL PROPERTIES OF AQUEOUS SOLUTIONS: Electrolytic Properties Ionic Compounds in Water,
Molecular Compounds in Water, Strong and Weak Electrolytes PRECIPITATION REACTIONS: Solubility
Guidelines for Ionic Compounds, Exchange
• (Metathesis) Reactions, Ionic Equations
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ACIDS, BASES, AND NEUTRALIZATION REACTIONS: Acids Bases Strong and Weak Acids and Bases,
Identifying Strong and Weak Electrolytes, Neutralization, Reactions and Salts, Neutralization Reactions with Gas
Formation
• OXIDATION-REDUCTION REACTIONS: Oxidation and Reduction, Oxidation Numbers, Oxidation of Metals
by Acids and Salts, The Activity Series
• CONCENTRATIONS OF SOLUTIONS: Molarity, Expressing the Concentration of an Electrolyte,
Interconverting Molarity, Moles, and Volume Dilution
• SOLUTION STOICHIOMETRY AND CHEMICAL ANALYSIS: Titrations
5. Thermochemistry
• THE NATURE OF ENERGY: Kinetic Energy and Potential Energy, Units of Energy, System and Surroundings,
Transferring Energy, Work and Heat
• THE FIRST LAW OF THERMODYNAMICS: Internal Energy, Relating to Heat and Work, Endothermic and
Exothermic Processes, State Functions,
• ENTHALPY
• ENTHALPIES OF REACTION
• CALORIMETRY: Heat Capacity and Specific Heat, Constant-Pressure Calorimetry, Bomb Calorimetry
(Constant-Volume Calorimetry)
• HESS’S LAW
• ENTHALPIES OF FORMATION: Using Enthalpies of Formation to Calculate Enthalpies of Reaction
• FOODS AND FUELS: Foods, Fuels, Other Energy Sources
6. Electronic Structure of Atoms
• THE W A VE NA TURE OF LIGHT
• QUANTIZED ENERGY AND PHOTONS: Hot Objects and the Quantization of Energy, The Photoelectric Effect
and Photons
• LINE SPECTRA AND THE BOHR MODEL: Line Spectra, Bohr’s Model, The Energy States of the Hydrogen
Atom, Limitations of the Bohr Model
• THE W A VE BEHA VIOR OF MA TTER: The Uncertainty Principle
• QUANTUM MECHANICS AND ATOMIC ORBITALS: Orbitals and Quantum Numbers
• REPRESENTATIONS OF ORBITALS: The s Orbitals, The p Orbitals, The d and f Orbitals
• MANY-ELECTRON ATOMS: Orbitals and Their Energies, Electron Spin and the Pauli Exclusion Principle
• ELECTRON CONFIGURATIONS: Hund’s Rule, Condensed Electron Configurations, Transition Metals, The
Lanthanides and Actinides
• ELECTRON CONFIGURATIONS AND THE PERIODIC TABLE: Anomalous Electron Configurations
7. Periodic Properties of the Elements
• DEVELOPMENT OF THE PERIODIC TABLE
• EFFECTIVE NUCLEAR CHARGE
• SIZES OF ATOMS AND IONS: Periodic Trends in Atomic Radii, Periodic Trends in Ionic Radii
• IONIZATION ENERGY: Variations in Successive Ionization Energies, Periodic Trends in First Ionization
Energies, Electron Configurations of Ions
• ELECTRON AFFINITIES
• METALS, NONMETALS, AND METALLOIDS: Metals, Nonmetals, Metalloids
• TRENDS FOR GROUP A AND GROUP A METALS: Group A - The Alkali Metals, Group A - The Alkaline
Earth Metals
• TRENDS FOR SELECTED NONMETALS: Hydrogen, Group A - The Oxygen Group, Group A - The
• Halogens, Group A - The Noble Gases
8. Basic Concepts of Chemical Bonding
• LEWIS SYMBOLS AND THE OCTET RULE: The Octet Rule
• IONIC BONDING: Energetics of Ionic Bond Formation, Electron Configurations of Ions, of the s- and p-Block
Elements, Transition-Metal Ions
• COV ALENT BONDING: Lewis Structures Multiple Bonds
• BOND POLARITY AND ELECTRONEGATIVITY: Electronegativity, Electronegativity and Bond Polarity,
Dipole Moments, Differentiating Ionic and Covalent Bonding
• DRAWING LEWIS STRUCTURES: Formal Charge and Alternative Lewis Structures
• RESONANCE STRUCTURES: Resonance in Benzene
• EXCEPTIONS TO THE OCTET RULE: Odd Number of Electrons, Less than an Octet of Valence Electrons,
More than an Octet of Valence Electrons
• STRENGTHS OF COVALENT BONDS: Bond Enthalpies and the Enthalpies of Reactions, Bond Enthalpy and
•
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Bond Length
9. Molecular Geometry and Bonding Theories
• MOLECULAR SHAPES
• THE VSEPR MODEL: Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles, Molecules with
Expanded Valence Shells, Shapes of Larger Molecules
• MOLECULAR SHAPE AND MOLECULAR POLARITY
• COVALENT BONDING AND ORBITAL OVERLAP
• HYBRID ORBITALS: sp Hybrid Orbitals, sp2 and sp3 Hybrid Orbitals, Hybrid Orbital Summary
• MULTIPLE BONDS: Resonance Structures, Delocalization, and Bonding
Problems
1. (a) After the label fell off a bottle containing a clear liquid believed to be benzene, a chemist measured the density of the
liquid to verify its identity. A 25.0-mL portion of the liquid had a mass of 21.95 g. A chemistry handbook lists the
density of benzene at 15 °C as 0.8787 g/mL. Is the calculated density in agreement with the tabulated value? (b) An
experiment requires 15.0 g of cyclohexane, whose density at 25 °C is 0.7781 g/mL. What volume of cyclohexane should
be used? (c) A spherical ball of lead has a diameter of 5.0 cm. What is the mass of the sphere if lead has a density of
11.34 g.cm3? (The volume of a sphere is (4/3)πr3where r is the radius.)
2. Gold is alloyed (mixed) with other metals to increase its hardness in making jewelry. (a) Consider a piece of gold
jewelry that weighs 9.85 g and has a volume of 0.675 cm3. The jewelry contains only gold and silver, which have
densities of 19.3 g/cm3 and 10.5 g/cm3, respectively. If the total volume of the jewelry is the sum of the volumes of the
gold and silver that it contains, calculate the percentage of gold (by mass) in the jewelry. (b) The relative amount of gold
in an alloy is commonly expressed in units of carats. Pure gold is 24 carat, and the percentage of gold in an alloy is
given as a percentage of this value. For example, an alloy that is 50% gold is 12 carat. State the purity of the gold
jewelry in carats.
3. The following diagram is a representation of 20 atoms of a fictitious element, which we will call nevadium (Nv). The
red spheres are 293Nv, and the blue spheres are 295Nv. (a) Assuming that this sample is a statistically representative
sample of the element, calculate the percent abundance of each element. (b) If the mass of 293Nv is 293.15 amu and that
of 295Nv is 295.15 amu, what is the atomic weight of Nv? [Section 2.4]
4. Name the following ionic compounds: (a) Li2O, (b) FeCl3 , (c) NaClO, (d) CaSO3 , (e) Cu(OH)2, (f) Fe(NO3)2 (g)
Ca(CH3COO)2, (h) Cr2(CO3)3 , (i) K2CrO4 , (j) (NH4)2SO4 .
5. Balance the following equations:
(a) Li (s) + N2 (g) à Li3N (s)
(b) TiCl4 (l) + H2O (l) à TiO2 (s) + HCl (aq)
(c) NH4NO3 (s) à N2 (g) + O2 (g) + H2O (g)
(d) Ca3P2 (s) + H2O (l) à Ca(OH)2 (aq) + PH3 (g)
(e) Al(OH)3 (s) + H2SO4 (aq) à Al2(SO4)3 (aq) + H2O (l)
(f) AgNO3 (aq) + Na2CO3 (aq) à Ag2NO3 (s) + Na2CO3 (aq)
(g) C2H5NH2 (g) + O2 (g) à CO2 (g) + H2O (g) + N2 (g)
6. Determine the formula weights of each of the following compounds: (a) nitric acid, HN3 ; (b) KMnO4; (c) Ca3(PO4)2; (d)
quartz, SiO2; (e) gallium sulfide, (f) chromium(III) sulfate, (g) phosphorus trichloride.
7. Calculate the following quantities:
(a) mass, in grams, of 1.50 x 10-2 mol of CdS
(b) number of moles of NH4Cl in 86.6 g of this substance
(c) number of molecules in 8.447 x 10-2 mol C6H6
(d) number of O atoms in 6.25 x 10-3 mol Al(NO3)3
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8. (a) Combustion analysis of toluene, a common organic solvent, gives 5.86 mg of CO2 and 1.37 mg of H2O. If the
compound contains only carbon and hydrogen, what is its empirical formula? (b) Menthol, the substance we can smell
in mentholated cough drops, is composed of C, H, and O. A 0.1005-g sample of menthol is combusted, producing
0.2829 g of CO2 and 0.1159 g of H2O. What is the empirical formula for menthol? If menthol has a molar mass of 156
g/mol, what is its molecular formula?
9. The complete combustion of octane, C8H18, the main component of gasoline, proceeds as follows:
2 C8H18 (l) + 25 O2 (g) à 16 CO2 (g) + 18 H2O (g)
(a) How many moles of O2 are needed to burn 1.50 mol of C8H18?
(b) How many grams of O2 are needed to burn 10.0 g of C8H18?
(c) Octane has a density of 0.692 g/mL at 20 °C. How many grams of O2 are required to burn 15.0 gal of C8H18 (the
capacity of an average fuel tank)?
(d) How many grams of CO2 are produced when 15.0 gal of C8H18 are combusted?
10. Write balanced net ionic equations for the reactions that occur in each of the following cases. Identify the spectator ion
or ions in each reaction.
(a) Cr2(SO4)3 (aq) + (NH4)2CO3 (aq) à
(b) Ba(NO3)2 (aq) + K2SO4 (aq) à
(c) Fe(NO3)2 (aq) + KOH (aq) à
11. Classify each of the following substances as a nonelectrolyte, weak electrolyte, or strong electrolyte in water: (a)
H2SO3, (b) C2H5OH (ethanol), (c) NH3, (d) KClO3, (e) Cu(NO3)2.
12. (a) Starting with solid sucrose, C12H22O11, describe how you would prepare 250 mL of a 0.250 M sucrose solution. (b)
Describe how you would prepare 350.0 mL of 0.100 M C12H22O11 starting with 3.00 L of 1.50 M C12H22O11.
13. You choose to investigate some of the solubility guidelines for two ions not listed in Table 4.1, the chromate ion (CrO42) and the oxalate ion (C2O42- ). You are given 0.01 M solutions (A, B, C, D) of four water-soluble salts:
Solution
Solute
Color of Solution
A
Na2CrO4
Yellow
B
(NH4)2C2O4
Colorless
C
AgNO3
Colorless
D
CaCl2
Colorless
When these solutions are mixed, the following observations are made:
Expt
Solutions
Number
Mixed
Result
1
A+B
No precipitate, yellow solution
2
A+C
Red precipitate forms
3
A+D
Yellow precipitate forms
4
B+C
White precipitate forms
5
B+D
White precipitate forms
6
C+D
White precipitate forms
(a) Write a net ionic equation for the reaction that occurs in each of the experiments. (b) Identify the precipitate formed,
if any, in each of the experiments.
14. Calculate ΔE and determine whether the process is endothermic or exothermic for the following cases: (a) q = 0.763 kJ
and w = –840 J; (b) a system releases 66.1 kJ of heat to its surroundings while the surroundings do 44.0 kJ of work on
the system; (c) the system absorbs 7.25 kJ of heat from the surroundings while its volume remains constant (assume that
only P-V work can be done).
15. Using values from any table of enthalpies, calculate the value of ΔH° for each of the following reactions:
(a) CaO (s) 2 HCl (g) à CaCl2 (s) + H2O (g)
(b) 4 FeO (s) + O2 (g) à 2 Fe2O3 (s)
(c) 2 CuO (s) + NO (g) à Cu2O (s) + NO2 (g)
(d) 4 NH3 (g) + O2 (g) à 2 N2H4 (g) + 2 H2O (l)
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16. (a) Calculate the energy of a photon of electromagnetic radiation whose frequency is 6.74 x 1012 s-1. (b) Calculate the
energy of a photon of radiation whose wavelength is 322 nm. (c) What wavelength of radiation has photons of energy
2.87 x 10-18 J.
17. Bohr’s model can be used for hydrogen-like ions—ions that have only one electron, such as He+ and Li2+. (a) Why is the
Bohr model applicable to He+ ions but not to neutral He atoms? (b) The ground-state energies of H, He+, and Li2+ are
tabulated as follows:
Atom or ion
Ground-state energy
H
-2.18 x 10-18 J
He+
-8.72 x 10-18 J
Li2+
-1.96 x 10-17 J
By examining these numbers, propose a relationship between the ground-state energy of hydrogen-like systems and the
nuclear charge, Z. (c) Use the relationship you derive in part (b) to predict the ground-state energy of the C5+ ion.
18. Write the Lewis symbol for atoms of each of the following elements: (a) Al, (b) Br, (c) Ar, (d) Sr.
19. Which of the following bonds are polar: (a) B – F, (b) Cl – Cl, (c) Se – O, (d) H – I? Which is the more electronegative
atom in each polar bond?
20. Draw the Lewis structures for each of the following ions or molecules. Identify those that do not obey the octet rule,
and explain why they do not: (a) SO32-, (b) AlH3, (c) N3-, (d) CH2Cl2, (e) SbF5.
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