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Introductory Chemistry - 1st Canadian Edition: End-of-Chapter Material

Introductory Chemistry - 1st Canadian Edition
End-of-Chapter Material
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table of contents
  1. Cover
  2. Title Page
  3. Copyright
  4. Table Of Contents
  5. Acknowledgments
  6. Dedication
  7. About BCcampus Open Education
  8. Chapter 1. What is Chemistry
    1. Some Basic Definitions
    2. Chemistry as a Science
  9. Chapter 2. Measurements
    1. Expressing Numbers
    2. Significant Figures
    3. Converting Units
    4. Other Units: Temperature and Density
    5. Expressing Units
    6. End-of-Chapter Material
  10. Chapter 3. Atoms, Molecules, and Ions
    1. Acids
    2. Ions and Ionic Compounds
    3. Masses of Atoms and Molecules
    4. Molecules and Chemical Nomenclature
    5. Atomic Theory
    6. End-of-Chapter Material
  11. Chapter 4. Chemical Reactions and Equations
    1. The Chemical Equation
    2. Types of Chemical Reactions: Single- and Double-Displacement Reactions
    3. Ionic Equations: A Closer Look
    4. Composition, Decomposition, and Combustion Reactions
    5. Oxidation-Reduction Reactions
    6. Neutralization Reactions
    7. End-of-Chapter Material
  12. Chapter 5. Stoichiometry and the Mole
    1. Stoichiometry
    2. The Mole
    3. Mole-Mass and Mass-Mass Calculations
    4. Limiting Reagents
    5. The Mole in Chemical Reactions
    6. Yields
    7. End-of-Chapter Material
  13. Chapter 6. Gases
    1. Pressure
    2. Gas Laws
    3. Other Gas Laws
    4. The Ideal Gas Law and Some Applications
    5. Gas Mixtures
    6. Kinetic Molecular Theory of Gases
    7. Molecular Effusion and Diffusion
    8. Real Gases
    9. End-of-Chapter Material
  14. Chapter 7. Energy and Chemistry
    1. Formation Reactions
    2. Energy
    3. Stoichiometry Calculations Using Enthalpy
    4. Enthalpy and Chemical Reactions
    5. Work and Heat
    6. Hess’s Law
    7. End-of-Chapter Material
  15. Chapter 8. Electronic Structure
    1. Light
    2. Quantum Numbers for Electrons
    3. Organization of Electrons in Atoms
    4. Electronic Structure and the Periodic Table
    5. Periodic Trends
    6. End-of-Chapter Material
  16. Chapter 9. Chemical Bonds
    1. Lewis Electron Dot Diagrams
    2. Electron Transfer: Ionic Bonds
    3. Covalent Bonds
    4. Other Aspects of Covalent Bonds
    5. Violations of the Octet Rule
    6. Molecular Shapes and Polarity
    7. Valence Bond Theory and Hybrid Orbitals
    8. Molecular Orbitals
    9. End-of-Chapter Material
  17. Chapter 10. Solids and Liquids
    1. Properties of Liquids
    2. Solids
    3. Phase Transitions: Melting, Boiling, and Subliming
    4. Intermolecular Forces
    5. End-of-Chapter Material
  18. Chapter 11. Solutions
    1. Colligative Properties of Solutions
    2. Concentrations as Conversion Factors
    3. Quantitative Units of Concentration
    4. Colligative Properties of Ionic Solutes
    5. Some Definitions
    6. Dilutions and Concentrations
    7. End-of-Chapter Material
  19. Chapter 12. Acids and Bases
    1. Acid-Base Titrations
    2. Strong and Weak Acids and Bases and Their Salts
    3. Brønsted-Lowry Acids and Bases
    4. Arrhenius Acids and Bases
    5. Autoionization of Water
    6. Buffers
    7. The pH Scale
    8. End-of-Chapter Material
  20. Chapter 13. Chemical Equilibrium
    1. Chemical Equilibrium
    2. The Equilibrium Constant
    3. Shifting Equilibria: Le Chatelier’s Principle
    4. Calculating Equilibrium Constant Values
    5. Some Special Types of Equilibria
    6. End-of-Chapter Material
  21. Chapter 14. Oxidation and Reduction
    1. Oxidation-Reduction Reactions
    2. Balancing Redox Reactions
    3. Applications of Redox Reactions: Voltaic Cells
    4. Electrolysis
    5. End-of-Chapter Material
  22. Chapter 15. Nuclear Chemistry
    1. Units of Radioactivity
    2. Uses of Radioactive Isotopes
    3. Half-Life
    4. Radioactivity
    5. Nuclear Energy
    6. End-of-Chapter Material
  23. Chapter 16. Organic Chemistry
    1. Hydrocarbons
    2. Branched Hydrocarbons
    3. Alkyl Halides and Alcohols
    4. Other Oxygen-Containing Functional Groups
    5. Other Functional Groups
    6. Polymers
    7. End-of-Chapter Material
  24. Chapter 17. Kinetics
    1. Factors that Affect the Rate of Reactions
    2. Reaction Rates
    3. Rate Laws
    4. Concentration–Time Relationships: Integrated Rate Laws
    5. Activation Energy and the Arrhenius Equation
    6. Reaction Mechanisms
    7. Catalysis
    8. End-of-Chapter Material
  25. Chapter 18. Chemical Thermodynamics
    1. Spontaneous Change
    2. Entropy and the Second Law of Thermodynamics
    3. Measuring Entropy and Entropy Changes
    4. Gibbs Free Energy
    5. Spontaneity: Free Energy and Temperature
    6. Free Energy under Nonstandard Conditions
    7. End-of-Chapter Material
  26. Appendix A: Periodic Table of the Elements
  27. Appendix B: Selected Acid Dissociation Constants at 25°C
  28. Appendix C: Solubility Constants for Compounds at 25°C
  29. Appendix D: Standard Thermodynamic Quantities for Chemical Substances at 25°C
  30. Appendix E: Standard Reduction Potentials by Value
  31. Glossary
  32. About the Authors
  33. Versioning History

End-of-Chapter Material

Additional Exercises

  1. What is the work when 124 mL of gas contract to 72.0 mL under an external pressure of 822 torr?
  2. What is the work when 2,345 mL of gas contract to 887 mL under an external pressure of 348 torr?
  3. A 3.77 L volume of gas is exposed to an external pressure of 1.67 atm. As the gas contracts, 156 J of work are added to the gas. What is the final volume of the gas?
  4. A 457 mL volume of gas contracts when 773 torr of external pressure act on it. If 27.4 J of work are added to the gas, what is its final volume?
  5. What is the heat when 1,744 g of Hg increase in temperature by 334°C? Express your final answer in kJ.
  6. What is the heat when 13.66 kg of Fe cool by 622°C? Express your final answer in kJ.
  7. What is final temperature when a 45.6 g sample of Al at 87.3°C gains 188 J of heat?
  8. What is final temperature when 967 g of Au at 557°C lose 559 J of heat?
  9. Plants take CO2 and H2O and make glucose (C6H12O6) and O2. Write a balanced thermochemical equation for this process. Use data in Table 7.1 “Enthalpies of Formation for Various Substances.”
  10. Exercise 9 described the formation of glucose in plants, which take in CO2 and H2O and give off O2. Is this process exothermic or endothermic? If exothermic, where does the energy go? If endothermic, where does the energy come from?
  11. The basic reaction in the refining of aluminum is to take Al2O3(s) and turn it into Al(s) and O2(g). Write the balanced thermochemical equation for this process. Use data in Table 7.1.
  12. Is the enthalpy change of the reaction H2O(ℓ) → H2O(g) zero or nonzero? Use data in Table 7.1 to determine the answer.
  13. What mass of H2O can be heated from 22°C to 80°C in the combustion of 1 mol of CH4? You will need the balanced thermochemical equation for the combustion of CH4. Use data in Table 7.1.
  14. What mass of H2O can be heated from 22°C to 80°C in the combustion of 1 mol of C2H6? You will need the balanced thermochemical equation for the combustion of C2H6. Use data in Table 7.1. Compare your answer to Exercise 13.
  15. What is the enthalpy change for the unknown reaction?

        \begin{align*} \ce{Pb(s)}+\ce{Cl2(g)}&\rightarrow\ce{PbCl2(s)} & \Delta H&=-359\text{ kJ} \\ \ce{PbCl2(s)}+\ce{Cl2(g)}&\rightarrow \ce{PbCl4(\ell)} & \Delta H&=? \\ \ce{Pb(s)}+\ce{2Cl2(g)}&\rightarrow \ce{PbCl4(\ell)} & \Delta H&=-329\text{ kJ} \end{align*}

  16. What is the enthalpy change for the unknown reaction?

        \begin{align*} \ce{P(s)}+\ce{\dfrac{3}{2}Br2(\ell)}&\rightarrow \ce{PBr3(\ell)} & \Delta H&=-185\text{ kJ} \\ \ce{PI3(s)}&\rightarrow \ce{P(s)}+\ce{\dfrac{3}{2}I2(s)} & \Delta H&=? \\ \ce{PI3}+\ce{\dfrac{3}{2}Br2(\ell)}&\rightarrow \ce{PBr3(\ell)}+\ce{\dfrac{3}{2}I2(s)} & \Delta H&=-139\text{ kJ} \end{align*}

  17. What is the ΔH for the reaction C(s, gra) → C(s, dia)? The label gra means graphite, and the label dia means diamond. What does your answer mean?
  18. Without consulting any tables, determine the ΔH for the reaction H2O(ℓ, 25°C) → H2O(ℓ, 25°C). Explain your answer.

Answers

  1. 5.70 J
  1. 4.69 L
  1. 80.97 kJ
  1. 91.9°C
  1. \ce{6CO2(g)}+\ce{6H2O(\ell)}\rightarrow \ce{C6H12O6(s)}+\ce{6O2(g)}\quad\Delta H=2,799\text{ kJ}
  1. \ce{2Al2O3(s)}\rightarrow \ce{4Al(s)}+\ce{3O2(g)}\quad\Delta H=3,351.4\text{ kJ}
  1. 3,668 g
  1. ΔH = 30 kJ
  1. ΔH = 1.897 kJ; the reaction is endothermic.

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Copyright © 2014

                                by Jessie A. Key

            Introductory Chemistry - 1st Canadian Edition by Jessie A. Key is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License, except where otherwise noted.
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