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Chemistry 2e

Exercises

Chemistry 2eExercises
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Table of contents
  1. Preface
  2. 1 Essential Ideas
    1. Introduction
    2. 1.1 Chemistry in Context
    3. 1.2 Phases and Classification of Matter
    4. 1.3 Physical and Chemical Properties
    5. 1.4 Measurements
    6. 1.5 Measurement Uncertainty, Accuracy, and Precision
    7. 1.6 Mathematical Treatment of Measurement Results
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  3. 2 Atoms, Molecules, and Ions
    1. Introduction
    2. 2.1 Early Ideas in Atomic Theory
    3. 2.2 Evolution of Atomic Theory
    4. 2.3 Atomic Structure and Symbolism
    5. 2.4 Chemical Formulas
    6. 2.5 The Periodic Table
    7. 2.6 Ionic and Molecular Compounds
    8. 2.7 Chemical Nomenclature
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  4. 3 Composition of Substances and Solutions
    1. Introduction
    2. 3.1 Formula Mass and the Mole Concept
    3. 3.2 Determining Empirical and Molecular Formulas
    4. 3.3 Molarity
    5. 3.4 Other Units for Solution Concentrations
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  5. 4 Stoichiometry of Chemical Reactions
    1. Introduction
    2. 4.1 Writing and Balancing Chemical Equations
    3. 4.2 Classifying Chemical Reactions
    4. 4.3 Reaction Stoichiometry
    5. 4.4 Reaction Yields
    6. 4.5 Quantitative Chemical Analysis
    7. Key Terms
    8. Key Equations
    9. Summary
    10. Exercises
  6. 5 Thermochemistry
    1. Introduction
    2. 5.1 Energy Basics
    3. 5.2 Calorimetry
    4. 5.3 Enthalpy
    5. Key Terms
    6. Key Equations
    7. Summary
    8. Exercises
  7. 6 Electronic Structure and Periodic Properties of Elements
    1. Introduction
    2. 6.1 Electromagnetic Energy
    3. 6.2 The Bohr Model
    4. 6.3 Development of Quantum Theory
    5. 6.4 Electronic Structure of Atoms (Electron Configurations)
    6. 6.5 Periodic Variations in Element Properties
    7. Key Terms
    8. Key Equations
    9. Summary
    10. Exercises
  8. 7 Chemical Bonding and Molecular Geometry
    1. Introduction
    2. 7.1 Ionic Bonding
    3. 7.2 Covalent Bonding
    4. 7.3 Lewis Symbols and Structures
    5. 7.4 Formal Charges and Resonance
    6. 7.5 Strengths of Ionic and Covalent Bonds
    7. 7.6 Molecular Structure and Polarity
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  9. 8 Advanced Theories of Covalent Bonding
    1. Introduction
    2. 8.1 Valence Bond Theory
    3. 8.2 Hybrid Atomic Orbitals
    4. 8.3 Multiple Bonds
    5. 8.4 Molecular Orbital Theory
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  10. 9 Gases
    1. Introduction
    2. 9.1 Gas Pressure
    3. 9.2 Relating Pressure, Volume, Amount, and Temperature: The Ideal Gas Law
    4. 9.3 Stoichiometry of Gaseous Substances, Mixtures, and Reactions
    5. 9.4 Effusion and Diffusion of Gases
    6. 9.5 The Kinetic-Molecular Theory
    7. 9.6 Non-Ideal Gas Behavior
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  11. 10 Liquids and Solids
    1. Introduction
    2. 10.1 Intermolecular Forces
    3. 10.2 Properties of Liquids
    4. 10.3 Phase Transitions
    5. 10.4 Phase Diagrams
    6. 10.5 The Solid State of Matter
    7. 10.6 Lattice Structures in Crystalline Solids
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  12. 11 Solutions and Colloids
    1. Introduction
    2. 11.1 The Dissolution Process
    3. 11.2 Electrolytes
    4. 11.3 Solubility
    5. 11.4 Colligative Properties
    6. 11.5 Colloids
    7. Key Terms
    8. Key Equations
    9. Summary
    10. Exercises
  13. 12 Kinetics
    1. Introduction
    2. 12.1 Chemical Reaction Rates
    3. 12.2 Factors Affecting Reaction Rates
    4. 12.3 Rate Laws
    5. 12.4 Integrated Rate Laws
    6. 12.5 Collision Theory
    7. 12.6 Reaction Mechanisms
    8. 12.7 Catalysis
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  14. 13 Fundamental Equilibrium Concepts
    1. Introduction
    2. 13.1 Chemical Equilibria
    3. 13.2 Equilibrium Constants
    4. 13.3 Shifting Equilibria: Le Châtelier’s Principle
    5. 13.4 Equilibrium Calculations
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  15. 14 Acid-Base Equilibria
    1. Introduction
    2. 14.1 Brønsted-Lowry Acids and Bases
    3. 14.2 pH and pOH
    4. 14.3 Relative Strengths of Acids and Bases
    5. 14.4 Hydrolysis of Salts
    6. 14.5 Polyprotic Acids
    7. 14.6 Buffers
    8. 14.7 Acid-Base Titrations
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  16. 15 Equilibria of Other Reaction Classes
    1. Introduction
    2. 15.1 Precipitation and Dissolution
    3. 15.2 Lewis Acids and Bases
    4. 15.3 Coupled Equilibria
    5. Key Terms
    6. Key Equations
    7. Summary
    8. Exercises
  17. 16 Thermodynamics
    1. Introduction
    2. 16.1 Spontaneity
    3. 16.2 Entropy
    4. 16.3 The Second and Third Laws of Thermodynamics
    5. 16.4 Free Energy
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  18. 17 Electrochemistry
    1. Introduction
    2. 17.1 Review of Redox Chemistry
    3. 17.2 Galvanic Cells
    4. 17.3 Electrode and Cell Potentials
    5. 17.4 Potential, Free Energy, and Equilibrium
    6. 17.5 Batteries and Fuel Cells
    7. 17.6 Corrosion
    8. 17.7 Electrolysis
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  19. 18 Representative Metals, Metalloids, and Nonmetals
    1. Introduction
    2. 18.1 Periodicity
    3. 18.2 Occurrence and Preparation of the Representative Metals
    4. 18.3 Structure and General Properties of the Metalloids
    5. 18.4 Structure and General Properties of the Nonmetals
    6. 18.5 Occurrence, Preparation, and Compounds of Hydrogen
    7. 18.6 Occurrence, Preparation, and Properties of Carbonates
    8. 18.7 Occurrence, Preparation, and Properties of Nitrogen
    9. 18.8 Occurrence, Preparation, and Properties of Phosphorus
    10. 18.9 Occurrence, Preparation, and Compounds of Oxygen
    11. 18.10 Occurrence, Preparation, and Properties of Sulfur
    12. 18.11 Occurrence, Preparation, and Properties of Halogens
    13. 18.12 Occurrence, Preparation, and Properties of the Noble Gases
    14. Key Terms
    15. Summary
    16. Exercises
  20. 19 Transition Metals and Coordination Chemistry
    1. Introduction
    2. 19.1 Occurrence, Preparation, and Properties of Transition Metals and Their Compounds
    3. 19.2 Coordination Chemistry of Transition Metals
    4. 19.3 Spectroscopic and Magnetic Properties of Coordination Compounds
    5. Key Terms
    6. Summary
    7. Exercises
  21. 20 Organic Chemistry
    1. Introduction
    2. 20.1 Hydrocarbons
    3. 20.2 Alcohols and Ethers
    4. 20.3 Aldehydes, Ketones, Carboxylic Acids, and Esters
    5. 20.4 Amines and Amides
    6. Key Terms
    7. Summary
    8. Exercises
  22. 21 Nuclear Chemistry
    1. Introduction
    2. 21.1 Nuclear Structure and Stability
    3. 21.2 Nuclear Equations
    4. 21.3 Radioactive Decay
    5. 21.4 Transmutation and Nuclear Energy
    6. 21.5 Uses of Radioisotopes
    7. 21.6 Biological Effects of Radiation
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  23. A | The Periodic Table
  24. B | Essential Mathematics
  25. C | Units and Conversion Factors
  26. D | Fundamental Physical Constants
  27. E | Water Properties
  28. F | Composition of Commercial Acids and Bases
  29. G | Standard Thermodynamic Properties for Selected Substances
  30. H | Ionization Constants of Weak Acids
  31. I | Ionization Constants of Weak Bases
  32. J | Solubility Products
  33. K | Formation Constants for Complex Ions
  34. L | Standard Electrode (Half-Cell) Potentials
  35. M | Half-Lives for Several Radioactive Isotopes
  36. Answer Key
    1. Chapter 1
    2. Chapter 2
    3. Chapter 3
    4. Chapter 4
    5. Chapter 5
    6. Chapter 6
    7. Chapter 7
    8. Chapter 8
    9. Chapter 9
    10. Chapter 10
    11. Chapter 11
    12. Chapter 12
    13. Chapter 13
    14. Chapter 14
    15. Chapter 15
    16. Chapter 16
    17. Chapter 17
    18. Chapter 18
    19. Chapter 19
    20. Chapter 20
    21. Chapter 21
  37. Index

16.1 Spontaneity

1.

What is a spontaneous reaction?

2.

What is a nonspontaneous reaction?

3.

Indicate whether the following processes are spontaneous or nonspontaneous.

(a) Liquid water freezing at a temperature below its freezing point

(b) Liquid water freezing at a temperature above its freezing point

(c) The combustion of gasoline

(d) A ball thrown into the air

(e) A raindrop falling to the ground

(f) Iron rusting in a moist atmosphere

4.

A helium-filled balloon spontaneously deflates overnight as He atoms diffuse through the wall of the balloon. Describe the redistribution of matter and/or energy that accompanies this process.

5.

Many plastic materials are organic polymers that contain carbon and hydrogen. The oxidation of these plastics in air to form carbon dioxide and water is a spontaneous process; however, plastic materials tend to persist in the environment. Explain.

16.2 Entropy

6.

In Figure 16.8 all possible distributions and microstates are shown for four different particles shared between two boxes. Determine the entropy change, ΔS, if the particles are initially evenly distributed between the two boxes, but upon redistribution all end up in Box (b).

7.

In Figure 16.8 all of the possible distributions and microstates are shown for four different particles shared between two boxes. Determine the entropy change, ΔS, for the system when it is converted from distribution (b) to distribution (d).

8.

How does the process described in the previous item relate to the system shown in Figure 16.4?

9.

Consider a system similar to the one in Figure 16.8, except that it contains six particles instead of four. What is the probability of having all the particles in only one of the two boxes in the case? Compare this with the similar probability for the system of four particles that we have derived to be equal to 18.18. What does this comparison tell us about even larger systems?

10.

Consider the system shown in Figure 16.9. What is the change in entropy for the process where the energy is initially associated only with particle A, but in the final state the energy is distributed between two different particles?

11.

Consider the system shown in Figure 16.9. What is the change in entropy for the process where the energy is initially associated with particles A and B, and the energy is distributed between two particles in different boxes (one in A-B, the other in C-D)?

12.

Arrange the following sets of systems in order of increasing entropy. Assume one mole of each substance and the same temperature for each member of a set.

(a) H2(g), HBrO4(g), HBr(g)

(b) H2O(l), H2O(g), H2O(s)

(c) He(g), Cl2(g), P4(g)

13.

At room temperature, the entropy of the halogens increases from I2 to Br2 to Cl2. Explain.

14.

Consider two processes: sublimation of I2(s) and melting of I2(s) (Note: the latter process can occur at the same temperature but somewhat higher pressure).

I 2 ( s ) ⟶ I 2 ( g ) I 2 ( s ) ⟶ I 2 ( g )

I 2 ( s ) ⟶ I 2 ( l ) I 2 ( s ) ⟶ I 2 ( l )

Is ΔS positive or negative in these processes? In which of the processes will the magnitude of the entropy change be greater?

15.

Indicate which substance in the given pairs has the higher entropy value. Explain your choices.

(a) C2H5OH(l) or C3H7OH(l)

(b) C2H5OH(l) or C2H5OH(g)

(c) 2H(g) or H(g)

16.

Predict the sign of the entropy change for the following processes.

(a) An ice cube is warmed to near its melting point.

(b) Exhaled breath forms fog on a cold morning.

(c) Snow melts.

17.

Predict the sign of the entropy change for the following processes. Give a reason for your prediction.

(a) Na+(aq)+Cl−(aq)⟶NaCl(s)Na+(aq)+Cl−(aq)⟶NaCl(s)

(b) 2Fe(s)+32O2(g)⟶Fe2O2(s)2Fe(s)+32O2(g)⟶Fe2O2(s)

(c) 2C6H14(l)+19O2(g)⟶14H2O(g)+12CO2(g)2C6H14(l)+19O2(g)⟶14H2O(g)+12CO2(g)

18.

Write the balanced chemical equation for the combustion of methane, CH4(g), to give carbon dioxide and water vapor. Explain why it is difficult to predict whether ΔS is positive or negative for this chemical reaction.

19.

Write the balanced chemical equation for the combustion of benzene, C6H6(l), to give carbon dioxide and water vapor. Would you expect ΔS to be positive or negative in this process?

16.3 The Second and Third Laws of Thermodynamics

20.

What is the difference between ΔS and ΔS° for a chemical change?

21.

Calculate ΔS°ΔS° for the following changes.

(a) SnCl4(l)⟶SnCl4(g)SnCl4(l)⟶SnCl4(g)

(b) CS2(g)⟶CS2(l)CS2(g)⟶CS2(l)

(c) Cu(s)⟶Cu(g)Cu(s)⟶Cu(g)

(d) H2O(l)⟶H2O(g)H2O(l)⟶H2O(g)

(e) 2H2(g)+O2(g)⟶2H2O(l)2H2(g)+O2(g)⟶2H2O(l)

(f) 2HCl(g)+Pb(s)⟶PbCl2(s)+H2(g)2HCl(g)+Pb(s)⟶PbCl2(s)+H2(g)

(g) Zn(s)+CuSO4(s)⟶Cu(s)+ZnSO4(s)Zn(s)+CuSO4(s)⟶Cu(s)+ZnSO4(s)

22.

Determine the entropy change for the combustion of liquid ethanol, C2H5OH, under the standard conditions to give gaseous carbon dioxide and liquid water.

23.

Determine the entropy change for the combustion of gaseous propane, C3H8, under the standard conditions to give gaseous carbon dioxide and water.

24.

“Thermite” reactions have been used for welding metal parts such as railway rails and in metal refining. One such thermite reaction is Fe2O3(s)+2Al(s)⟶Al2O3(s)+2Fe(s).Fe2O3(s)+2Al(s)⟶Al2O3(s)+2Fe(s). Is the reaction spontaneous at room temperature under standard conditions? During the reaction, the surroundings absorb 851.8 kJ/mol of heat.

25.

Using the relevant S°S° values listed in Appendix G, calculate ΔS°298ΔS°298 for the following changes:

(a) N2(g)+3H2(g)⟶2NH3(g)N2(g)+3H2(g)⟶2NH3(g)

(b) N2(g)+52O2(g)⟶N2O5(g)N2(g)+52O2(g)⟶N2O5(g)

26.

From the following information, determine ΔS°ΔS° for the following:

N ( g ) + O ( g ) ⟶ NO ( g ) Δ S ° = ? N ( g ) + O ( g ) ⟶ NO ( g ) Δ S ° = ?

N 2 ( g ) + O 2 ( g ) ⟶ 2 NO ( g ) Δ S ° = 24.8 J/K N 2 ( g ) + O 2 ( g ) ⟶ 2 NO ( g ) Δ S ° = 24.8 J/K

N 2 ( g ) ⟶ 2 N ( g ) Δ S ° = 115.0 J/K N 2 ( g ) ⟶ 2 N ( g ) Δ S ° = 115.0 J/K

O 2 ( g ) ⟶ 2 O ( g ) Δ S ° = 117.0 J/K O 2 ( g ) ⟶ 2 O ( g ) Δ S ° = 117.0 J/K

27.

By calculating ΔSuniv at each temperature, determine if the melting of 1 mole of NaCl(s) is spontaneous at 500 °C and at 700 °C.
 SNaCl(s)°=72.11Jmol·KSNaCl(l)°=95.06Jmol·KΔHfusion°=27.95 kJ/molSNaCl(s)°=72.11Jmol·KSNaCl(l)°=95.06Jmol·KΔHfusion°=27.95 kJ/mol

What assumptions are made about the thermodynamic information (entropy and enthalpy values) used to solve this problem?

28.

Use the standard entropy data in Appendix G to determine the change in entropy for each of the following reactions. All the processes occur at the standard conditions and 25 °C.

(a) MnO2(s)⟶Mn(s)+O2(g)MnO2(s)⟶Mn(s)+O2(g)

(b) H2(g)+Br2(l)⟶2HBr(g)H2(g)+Br2(l)⟶2HBr(g)

(c) Cu(s)+S(g)⟶CuS(s)Cu(s)+S(g)⟶CuS(s)

(d) 2LiOH(s)+CO2(g)⟶Li2CO3(s)+H2O(g)2LiOH(s)+CO2(g)⟶Li2CO3(s)+H2O(g)

(e) CH4(g)+O2(g)⟶C(s,graphite)+2H2O(g)CH4(g)+O2(g)⟶C(s,graphite)+2H2O(g)

(f) CS2(g)+3Cl2(g)⟶CCl4(g)+S2Cl2(g)CS2(g)+3Cl2(g)⟶CCl4(g)+S2Cl2(g)

29.

Use the standard entropy data in Appendix G to determine the change in entropy for each of the following reactions. All the processes occur at the standard conditions and 25 °C.

(a) C(s, graphite)+O2(g)⟶CO2(g)C(s, graphite)+O2(g)⟶CO2(g)

(b) O2(g)+N2(g)⟶2NO(g)O2(g)+N2(g)⟶2NO(g)

(c) 2Cu(s)+S(g)⟶Cu2S(s)2Cu(s)+S(g)⟶Cu2S(s)

(d) CaO(s)+H2O(l)⟶Ca(OH)2(s)CaO(s)+H2O(l)⟶Ca(OH)2(s)

(e) Fe2O3(s)+3CO(g)⟶2Fe(s)+3CO2(g)Fe2O3(s)+3CO(g)⟶2Fe(s)+3CO2(g)

(f) CaSO4·2H2O(s)⟶CaSO4(s)+2H2O(g)CaSO4·2H2O(s)⟶CaSO4(s)+2H2O(g)

16.4 Free Energy

30.

What is the difference between ΔG and ΔG° for a chemical change?

31.

A reaction has ΔH°ΔH° = 100 kJ/mol and ΔS°=250 J/mol·K.ΔS°=250 J/mol·K. Is the reaction spontaneous at room temperature? If not, under what temperature conditions will it become spontaneous?

32.

Explain what happens as a reaction starts with ΔG < 0 (negative) and reaches the point where ΔG = 0.

33.

Use the standard free energy of formation data in Appendix G to determine the free energy change for each of the following reactions, which are run under standard state conditions and 25 °C. Identify each as either spontaneous or nonspontaneous at these conditions.

(a) MnO2(s)⟶Mn(s)+O2(g)MnO2(s)⟶Mn(s)+O2(g)

(b) H2(g)+Br2(l)⟶2HBr(g)H2(g)+Br2(l)⟶2HBr(g)

(c) Cu(s)+S(g)⟶CuS(s)Cu(s)+S(g)⟶CuS(s)

(d) 2LiOH(s)+CO2(g)⟶Li2CO3(s)+H2O(g)2LiOH(s)+CO2(g)⟶Li2CO3(s)+H2O(g)

(e) CH4(g)+O2(g)⟶C(s,graphite)+2H2O(g)CH4(g)+O2(g)⟶C(s,graphite)+2H2O(g)

(f) CS2(g)+3Cl2(g)⟶CCl4(g)+S2Cl2(g)CS2(g)+3Cl2(g)⟶CCl4(g)+S2Cl2(g)

34.

Use the standard free energy data in Appendix G to determine the free energy change for each of the following reactions, which are run under standard state conditions and 25 °C. Identify each as either spontaneous or nonspontaneous at these conditions.

(a) C(s, graphite)+O2(g)⟶CO2(g)C(s, graphite)+O2(g)⟶CO2(g)

(b) O2(g)+N2(g)⟶2NO(g)O2(g)+N2(g)⟶2NO(g)

(c) 2Cu(s)+S(g)⟶Cu2S(s)2Cu(s)+S(g)⟶Cu2S(s)

(d) CaO(s)+H2O(l)⟶Ca(OH)2(s)CaO(s)+H2O(l)⟶Ca(OH)2(s)

(e) Fe2O3(s)+3CO(g)⟶2Fe(s)+3CO2(g)Fe2O3(s)+3CO(g)⟶2Fe(s)+3CO2(g)

(f) CaSO4·2H2O(s)⟶CaSO4(s)+2H2O(g)CaSO4·2H2O(s)⟶CaSO4(s)+2H2O(g)

35.

Given:
 P4(s)+5O2(g)⟶P4O10(s)ΔG°=−2697.0 kJ/molP4(s)+5O2(g)⟶P4O10(s)ΔG°=−2697.0 kJ/mol
 2H2(g)+O2(g)⟶2H2O(g)ΔG°=−457.18 kJ/mol2H2(g)+O2(g)⟶2H2O(g)ΔG°=−457.18 kJ/mol
 6H2O(g)+P4O10(s)⟶4H3PO4(l)ΔG°=−428.66 kJ/mol6H2O(g)+P4O10(s)⟶4H3PO4(l)ΔG°=−428.66 kJ/mol

(a) Determine the standard free energy of formation, ΔGf°,ΔGf°, for phosphoric acid.

(b) How does your calculated result compare to the value in Appendix G? Explain.

36.

Is the formation of ozone (O3(g)) from oxygen (O2(g)) spontaneous at room temperature under standard state conditions?

37.

Consider the decomposition of red mercury(II) oxide under standard state conditions.
 2HgO(s,red)⟶2Hg(l)+O2(g)2HgO(s,red)⟶2Hg(l)+O2(g)

(a) Is the decomposition spontaneous under standard state conditions?

(b) Above what temperature does the reaction become spontaneous?

38.

Among other things, an ideal fuel for the control thrusters of a space vehicle should decompose in a spontaneous exothermic reaction when exposed to the appropriate catalyst. Evaluate the following substances under standard state conditions as suitable candidates for fuels.

(a) Ammonia: 2NH3(g)⟶N2(g)+3H2(g)2NH3(g)⟶N2(g)+3H2(g)

(b) Diborane: B2H6(g)⟶2B(g)+3H2(g)B2H6(g)⟶2B(g)+3H2(g)

(c) Hydrazine: N2H4(g)⟶N2(g)+2H2(g)N2H4(g)⟶N2(g)+2H2(g)

(d) Hydrogen peroxide: H2O2(l)⟶H2O(g)+12O2(g)H2O2(l)⟶H2O(g)+12O2(g)

39.

Calculate ΔG° for each of the following reactions from the equilibrium constant at the temperature given.

(a) N2(g)+O2(g)⟶2NO(g)T=2000°CKp=4.1×10−4N2(g)+O2(g)⟶2NO(g)T=2000°CKp=4.1×10−4

(b) H2(g)+I2(g)⟶2HI(g)T=400°CKp=50.0H2(g)+I2(g)⟶2HI(g)T=400°CKp=50.0

(c) CO2(g)+H2(g)⟶CO(g)+H2O(g)T=980°CKp=1.67CO2(g)+H2(g)⟶CO(g)+H2O(g)T=980°CKp=1.67

(d) CaCO3(s)⟶CaO(s)+CO2(g)T=900°CKp=1.04CaCO3(s)⟶CaO(s)+CO2(g)T=900°CKp=1.04

(e) HF(aq)+H2O(l)⟶H3O+(aq)+F−(aq)T=25°CKp=7.2×10−4HF(aq)+H2O(l)⟶H3O+(aq)+F−(aq)T=25°CKp=7.2×10−4

(f) AgBr(s)⟶Ag+(aq)+Br−(aq)T=25°CKp=3.3×10−13AgBr(s)⟶Ag+(aq)+Br−(aq)T=25°CKp=3.3×10−13

40.

Calculate ΔG° for each of the following reactions from the equilibrium constant at the temperature given.

(a) Cl2(g)+Br2(g)⟶2BrCl(g)T=25°CKp=4.7×10−2Cl2(g)+Br2(g)⟶2BrCl(g)T=25°CKp=4.7×10−2

(b) 2SO2(g)+O2(g)⇌2SO3(g)T=500°CKp=48.22SO2(g)+O2(g)⇌2SO3(g)T=500°CKp=48.2

(c) H2O(l)⇌H2O(g)T=60°CKp=0.196H2O(l)⇌H2O(g)T=60°CKp=0.196

(d) CoO(s)+CO(g)⇌Co(s)+CO2(g)T=550°CKp=4.90×102CoO(s)+CO(g)⇌Co(s)+CO2(g)T=550°CKp=4.90×102

(e) CH3NH2(aq)+H2O(l)⟶CH3NH3+(aq)+OH−(aq)T=25°CKp=4.4×10−4CH3NH2(aq)+H2O(l)⟶CH3NH3+(aq)+OH−(aq)T=25°CKp=4.4×10−4

(f) PbI2(s)⟶Pb2+(aq)+2I−(aq)T=25°CKp=8.7×10−9PbI2(s)⟶Pb2+(aq)+2I−(aq)T=25°CKp=8.7×10−9

41.

Calculate the equilibrium constant at 25 °C for each of the following reactions from the value of ΔG° given.

(a) O2(g)+2F2(g)⟶2OF2(g)ΔG°=−9.2 kJO2(g)+2F2(g)⟶2OF2(g)ΔG°=−9.2 kJ

(b) I2(s)+Br2(l)⟶2IBr(g)ΔG°=7.3 kJI2(s)+Br2(l)⟶2IBr(g)ΔG°=7.3 kJ

(c) 2LiOH(s)+CO2(g)⟶Li2CO3(s)+H2O(g)ΔG°=−79 kJ2LiOH(s)+CO2(g)⟶Li2CO3(s)+H2O(g)ΔG°=−79 kJ

(d) N2O3(g)⟶NO(g)+NO2(g)ΔG°=−1.6 kJN2O3(g)⟶NO(g)+NO2(g)ΔG°=−1.6 kJ

(e) SnCl4(l)⟶SnCl4(l)ΔG°=8.0 kJSnCl4(l)⟶SnCl4(l)ΔG°=8.0 kJ

42.

Calculate the equilibrium constant at 25 °C for each of the following reactions from the value of ΔG° given.

(a) I2(s)+Cl2(g)⟶2ICl(g)ΔG°=−10.88 kJI2(s)+Cl2(g)⟶2ICl(g)ΔG°=−10.88 kJ

(b) H2(g)+I2(s)⟶2HI(g)ΔG°=3.4 kJH2(g)+I2(s)⟶2HI(g)ΔG°=3.4 kJ

(c) CS2(g)+3Cl2(g)⟶CCl4(g)+S2Cl2(g)ΔG°=−39 kJCS2(g)+3Cl2(g)⟶CCl4(g)+S2Cl2(g)ΔG°=−39 kJ

(d) 2SO2(g)+O2(g)⟶2SO3(g)ΔG°=−141.82 kJ2SO2(g)+O2(g)⟶2SO3(g)ΔG°=−141.82 kJ

(e) CS2(g)⟶CS2(l)ΔG°=−1.88 kJCS2(g)⟶CS2(l)ΔG°=−1.88 kJ

43.

Calculate the equilibrium constant at the temperature given.

(a) O2(g)+2F2(g)⟶2F2O(g)(T=100°C)O2(g)+2F2(g)⟶2F2O(g)(T=100°C)

(b) I2(s)+Br2(l)⟶2IBr(g)(T=0.0°C)I2(s)+Br2(l)⟶2IBr(g)(T=0.0°C)

(c) 2LiOH(s)+CO2(g)⟶Li2CO3(s)+H2O(g)(T=575°C)2LiOH(s)+CO2(g)⟶Li2CO3(s)+H2O(g)(T=575°C)

(d) N2O3(g)⟶NO(g)+NO2(g)(T=−10.0°C)N2O3(g)⟶NO(g)+NO2(g)(T=−10.0°C)

(e) SnCl4(l)⟶SnCl4(g)(T=200°C)SnCl4(l)⟶SnCl4(g)(T=200°C)

44.

Calculate the equilibrium constant at the temperature given.

(a) I2(s)+Cl2(g)⟶2ICl(g)(T=100°C)I2(s)+Cl2(g)⟶2ICl(g)(T=100°C)

(b) H2(g)+I2(s)⟶2HI(g)(T=0.0°C)H2(g)+I2(s)⟶2HI(g)(T=0.0°C)

(c) CS2(g)+3Cl2(g)⟶CCl4(g)+S2Cl2(g)(T=125°C)CS2(g)+3Cl2(g)⟶CCl4(g)+S2Cl2(g)(T=125°C)

(d) 2SO2(g)+O2(g)⟶2SO3(g)(T=675°C)2SO2(g)+O2(g)⟶2SO3(g)(T=675°C)

(e) CS2(g)⟶CS2(l)(T=90°C)CS2(g)⟶CS2(l)(T=90°C)

45.

Consider the following reaction at 298 K:
 N2O4(g)⇌2NO2(g)KP=0.142N2O4(g)⇌2NO2(g)KP=0.142

What is the standard free energy change at this temperature? Describe what happens to the initial system, where the reactants and products are in standard states, as it approaches equilibrium.

46.

Determine the normal boiling point (in kelvin) of dichloromethane, CH2Cl2. Find the actual boiling point using the Internet or some other source, and calculate the percent error in the temperature. Explain the differences, if any, between the two values.

47.

Under what conditions is N2O3(g)⟶NO(g)+NO2(g)N2O3(g)⟶NO(g)+NO2(g) spontaneous?

48.

At room temperature, the equilibrium constant (Kw) for the self-ionization of water is 1.00 ×× 10−14. Using this information, calculate the standard free energy change for the aqueous reaction of hydrogen ion with hydroxide ion to produce water. (Hint: The reaction is the reverse of the self-ionization reaction.)

49.

Hydrogen sulfide is a pollutant found in natural gas. Following its removal, it is converted to sulfur by the reaction 2H2S(g)+SO2(g)⇌38S8(s,rhombic)+2H2O(l).2H2S(g)+SO2(g)⇌38S8(s,rhombic)+2H2O(l). What is the equilibrium constant for this reaction? Is the reaction endothermic or exothermic?

50.

Consider the decomposition of CaCO3(s) into CaO(s) and CO2(g). What is the equilibrium partial pressure of CO2 at room temperature?

51.

In the laboratory, hydrogen chloride (HCl(g)) and ammonia (NH3(g)) often escape from bottles of their solutions and react to form the ammonium chloride (NH4Cl(s)), the white glaze often seen on glassware. Assuming that the number of moles of each gas that escapes into the room is the same, what is the maximum partial pressure of HCl and NH3 in the laboratory at room temperature? (Hint: The partial pressures will be equal and are at their maximum value when at equilibrium.)

52.

Benzene can be prepared from acetylene. 3C2H2(g)⇌C6H6(g).3C2H2(g)⇌C6H6(g). Determine the equilibrium constant at 25 °C and at 850 °C. Is the reaction spontaneous at either of these temperatures? Why is all acetylene not found as benzene?

53.

Carbon dioxide decomposes into CO and O2 at elevated temperatures. What is the equilibrium partial pressure of oxygen in a sample at 1000 °C for which the initial pressure of CO2 was 1.15 atm?

54.

Carbon tetrachloride, an important industrial solvent, is prepared by the chlorination of methane at 850 K.
 CH4(g)+4Cl2(g)⟶CCl4(g)+4HCl(g)CH4(g)+4Cl2(g)⟶CCl4(g)+4HCl(g)

What is the equilibrium constant for the reaction at 850 K? Would the reaction vessel need to be heated or cooled to keep the temperature of the reaction constant?

55.

Acetic acid, CH3CO2H, can form a dimer, (CH3CO2H)2, in the gas phase.
 2CH3CO2H(g)⟶(CH3CO2H)2(g)2CH3CO2H(g)⟶(CH3CO2H)2(g)

The dimer is held together by two hydrogen bonds with a total strength of 66.5 kJ per mole of dimer.

This Lewis structure shows a six-sided ring structure composed of a methyl group single bonded to a carbon, which is double bonded to an oxygen atom in an upward position and single bonded to an oxygen atom in a downward position. The lower oxygen is single bonded to a hydrogen, which is connected by a dotted line to an oxygen that is double bonded to a carbon in an upward position. This carbon is single bonded to a methyl group to its right and to an oxygen in the upward position that is single bonded to a hydrogen that is connected by a dotted line to the double bonded oxygen on the left.

At 25 °C, the equilibrium constant for the dimerization is 1.3 ×× 103 (pressure in atm). What is ΔS° for the reaction?

56.

Determine ΔGº for the following reactions.

(a) Antimony pentachloride decomposes at 448 °C. The reaction is:
 SbCl5(g)⟶SbCl3(g)+Cl2(g)SbCl5(g)⟶SbCl3(g)+Cl2(g)

An equilibrium mixture in a 5.00 L flask at 448 °C contains 3.85 g of SbCl5, 9.14 g of SbCl3, and 2.84 g of Cl2.

(b) Chlorine molecules dissociate according to this reaction:
 Cl2(g)⟶2Cl(g)Cl2(g)⟶2Cl(g)

1.00% of Cl2 molecules dissociate at 975 K and a pressure of 1.00 atm.

57.

Given that the ΔGf°ΔGf° for Pb2+(aq) and Cl−(aq) is −24.3 kJ/mole and −131.2 kJ/mole respectively, determine the solubility product, Ksp, for PbCl2(s).

58.

Determine the standard free energy change, ΔGf°,ΔGf°, for the formation of S2−(aq) given that the ΔGf°ΔGf° for Ag+(aq) and Ag2S(s) are 77.1 kJ/mole and −39.5 kJ/mole respectively, and the solubility product for Ag2S(s) is 8 ×× 10−51.

59.

Determine the standard enthalpy change, entropy change, and free energy change for the conversion of diamond to graphite. Discuss the spontaneity of the conversion with respect to the enthalpy and entropy changes. Explain why diamond spontaneously changing into graphite is not observed.

60.

The evaporation of one mole of water at 298 K has a standard free energy change of 8.58 kJ.
 H2O(l)⇌H2O(g)ΔG°=8.58 kJH2O(l)⇌H2O(g)ΔG°=8.58 kJ

(a) Is the evaporation of water under standard thermodynamic conditions spontaneous?

(b) Determine the equilibrium constant, KP, for this physical process.

(c) By calculating ∆G, determine if the evaporation of water at 298 K is spontaneous when the partial pressure of water, PH2O,PH2O, is 0.011 atm.

(d) If the evaporation of water were always nonspontaneous at room temperature, wet laundry would never dry when placed outside. In order for laundry to dry, what must be the value of PH2OPH2O in the air?

61.

In glycolysis, the reaction of glucose (Glu) to form glucose-6-phosphate (G6P) requires ATP to be present as described by the following equation:
 Glu+ATP⟶G6P+ADPΔG°=−17 kJGlu+ATP⟶G6P+ADPΔG°=−17 kJ

In this process, ATP becomes ADP summarized by the following equation:
 ATP⟶ADPΔG°=−30 kJATP⟶ADPΔG°=−30 kJ

Determine the standard free energy change for the following reaction, and explain why ATP is necessary to drive this process:
 Glu⟶G6PΔG°=?Glu⟶G6PΔG°=?

62.

One of the important reactions in the biochemical pathway glycolysis is the reaction of glucose-6-phosphate (G6P) to form fructose-6-phosphate (F6P):
 G6P⇌F6PΔG°=1.7 kJG6P⇌F6PΔG°=1.7 kJ

(a) Is the reaction spontaneous or nonspontaneous under standard thermodynamic conditions?

(b) Standard thermodynamic conditions imply the concentrations of G6P and F6P to be 1 M, however, in a typical cell, they are not even close to these values. Calculate ΔG when the concentrations of G6P and F6P are 120 μM and 28 μM respectively, and discuss the spontaneity of the forward reaction under these conditions. Assume the temperature is 37 °C.

63.

Without doing a numerical calculation, determine which of the following will reduce the free energy change for the reaction, that is, make it less positive or more negative, when the temperature is increased. Explain.

(a) N2(g)+3H2(g)⟶2NH3(g)N2(g)+3H2(g)⟶2NH3(g)

(b) HCl(g)+NH3(g)⟶NH4Cl(s)HCl(g)+NH3(g)⟶NH4Cl(s)

(c) (NH4)2Cr2O7(s)⟶Cr2O3(s)+4H2O(g)+N2(g)(NH4)2Cr2O7(s)⟶Cr2O3(s)+4H2O(g)+N2(g)

(d) 2Fe(s)+3O2(g)⟶Fe2O3(s)2Fe(s)+3O2(g)⟶Fe2O3(s)

64.

When ammonium chloride is added to water and stirred, it dissolves spontaneously and the resulting solution feels cold. Without doing any calculations, deduce the signs of ΔG, ΔH, and ΔS for this process, and justify your choices.

65.

An important source of copper is from the copper ore, chalcocite, a form of copper(I) sulfide. When heated, the Cu2S decomposes to form copper and sulfur described by the following equation:
 Cu2S(s)⟶Cu(s)+S(s)Cu2S(s)⟶Cu(s)+S(s)

(a) Determine ΔG°ΔG° for the decomposition of Cu2S(s).

(b) The reaction of sulfur with oxygen yields sulfur dioxide as the only product. Write an equation that describes this reaction, and determine ΔG°ΔG° for the process.

(c) The production of copper from chalcocite is performed by roasting the Cu2S in air to produce the Cu. By combining the equations from Parts (a) and (b), write the equation that describes the roasting of the chalcocite, and explain why coupling these reactions together makes for a more efficient process for the production of the copper.

66.

What happens to ΔGΔG (becomes more negative or more positive) for the following chemical reactions when the partial pressure of oxygen is increased?

(a) S(s)+O2(g)⟶SO2(g)S(s)+O2(g)⟶SO2(g)

(b) 2SO2(g)+O2(g)⟶2SO3(g)2SO2(g)+O2(g)⟶2SO3(g)

(c) HgO(s)⟶Hg(l)+O2(g)HgO(s)⟶Hg(l)+O2(g)

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