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Chemistry: Atoms First

Chapter 14

Chemistry: Atoms FirstChapter 14

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. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  4. 3 Electronic Structure and Periodic Properties of Elements
    1. Introduction
    2. 3.1 Electromagnetic Energy
    3. 3.2 The Bohr Model
    4. 3.3 Development of Quantum Theory
    5. 3.4 Electronic Structure of Atoms (Electron Configurations)
    6. 3.5 Periodic Variations in Element Properties
    7. 3.6 The Periodic Table
    8. 3.7 Molecular and Ionic Compounds
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  5. 4 Chemical Bonding and Molecular Geometry
    1. Introduction
    2. 4.1 Ionic Bonding
    3. 4.2 Covalent Bonding
    4. 4.3 Chemical Nomenclature
    5. 4.4 Lewis Symbols and Structures
    6. 4.5 Formal Charges and Resonance
    7. 4.6 Molecular Structure and Polarity
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  6. 5 Advanced Theories of Bonding
    1. Introduction
    2. 5.1 Valence Bond Theory
    3. 5.2 Hybrid Atomic Orbitals
    4. 5.3 Multiple Bonds
    5. 5.4 Molecular Orbital Theory
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  7. 6 Composition of Substances and Solutions
    1. Introduction
    2. 6.1 Formula Mass
    3. 6.2 Determining Empirical and Molecular Formulas
    4. 6.3 Molarity
    5. 6.4 Other Units for Solution Concentrations
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  8. 7 Stoichiometry of Chemical Reactions
    1. Introduction
    2. 7.1 Writing and Balancing Chemical Equations
    3. 7.2 Classifying Chemical Reactions
    4. 7.3 Reaction Stoichiometry
    5. 7.4 Reaction Yields
    6. 7.5 Quantitative Chemical Analysis
    7. Key Terms
    8. Key Equations
    9. Summary
    10. Exercises
  9. 8 Gases
    1. Introduction
    2. 8.1 Gas Pressure
    3. 8.2 Relating Pressure, Volume, Amount, and Temperature: The Ideal Gas Law
    4. 8.3 Stoichiometry of Gaseous Substances, Mixtures, and Reactions
    5. 8.4 Effusion and Diffusion of Gases
    6. 8.5 The Kinetic-Molecular Theory
    7. 8.6 Non-Ideal Gas Behavior
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  10. 9 Thermochemistry
    1. Introduction
    2. 9.1 Energy Basics
    3. 9.2 Calorimetry
    4. 9.3 Enthalpy
    5. 9.4 Strengths of Ionic and Covalent Bonds
    6. Key Terms
    7. Key Equations
    8. Summary
    9. 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 Thermodynamics
    1. Introduction
    2. 12.1 Spontaneity
    3. 12.2 Entropy
    4. 12.3 The Second and Third Laws of Thermodynamics
    5. 12.4 Free Energy
    6. Key Terms
    7. Key Equations
    8. Summary
    9. 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 Salt Solutions
    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 Multiple Equilibria
    5. Key Terms
    6. Key Equations
    7. Summary
    8. Exercises
  17. 16 Electrochemistry
    1. Introduction
    2. 16.1 Balancing Oxidation-Reduction Reactions
    3. 16.2 Galvanic Cells
    4. 16.3 Standard Reduction Potentials
    5. 16.4 The Nernst Equation
    6. 16.5 Batteries and Fuel Cells
    7. 16.6 Corrosion
    8. 16.7 Electrolysis
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  18. 17 Kinetics
    1. Introduction
    2. 17.1 Chemical Reaction Rates
    3. 17.2 Factors Affecting Reaction Rates
    4. 17.3 Rate Laws
    5. 17.4 Integrated Rate Laws
    6. 17.5 Collision Theory
    7. 17.6 Reaction Mechanisms
    8. 17.7 Catalysis
    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 Nuclear Chemistry
    1. Introduction
    2. 20.1 Nuclear Structure and Stability
    3. 20.2 Nuclear Equations
    4. 20.3 Radioactive Decay
    5. 20.4 Transmutation and Nuclear Energy
    6. 20.5 Uses of Radioisotopes
    7. 20.6 Biological Effects of Radiation
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  22. 21 Organic Chemistry
    1. Introduction
    2. 21.1 Hydrocarbons
    3. 21.2 Alcohols and Ethers
    4. 21.3 Aldehydes, Ketones, Carboxylic Acids, and Esters
    5. 21.4 Amines and Amides
    6. Key Terms
    7. Summary
    8. 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
1.

One example for NH3 as a conjugate acid: NH2+H+NH3;NH2+H+NH3; as a conjugate base: NH4+(aq)+OH(aq)NH3(aq)+H2O(l)NH4+(aq)+OH(aq)NH3(aq)+H2O(l)

3.

(a) H3O+(aq)H+(aq)+H2O(l);H3O+(aq)H+(aq)+H2O(l); (b) HCl(aq)H+(aq)+Cl(aq);HCl(aq)H+(aq)+Cl(aq); (c) NH3(aq)H+(aq)+NH2(aq);NH3(aq)H+(aq)+NH2(aq); (d) CH3CO2H(aq)H+(aq)+CH3CO2(aq);CH3CO2H(aq)H+(aq)+CH3CO2(aq); (e) NH4+(aq)H+(aq)+NH3(aq);NH4+(aq)H+(aq)+NH3(aq); (f) HSO4(aq)H+(aq)+SO42−(aq)HSO4(aq)H+(aq)+SO42−(aq)

5.

(a) H2O(l)+H+(aq)H3O+(aq);H2O(l)+H+(aq)H3O+(aq); (b) OH(aq)+H+(aq)H2O(l);OH(aq)+H+(aq)H2O(l); (c) NH3(aq)+H+(aq)NH4+(aq);NH3(aq)+H+(aq)NH4+(aq); (d) CN(aq)+H+(aq)HCN(aq);CN(aq)+H+(aq)HCN(aq); (e) S2−(aq)+H+(aq)HS(aq);S2−(aq)+H+(aq)HS(aq); (f) H2PO4(aq)+H+(aq)H3PO4(aq)H2PO4(aq)+H+(aq)H3PO4(aq)

7.

(a) H2O, O2−; (b) H3O+, OH; (c) H2CO3, CO32−;CO32−; (d) NH4+,NH4+, NH2;NH2; (e) H2SO4, SO42−;SO42−; (f) H3O2+,H3O2+, HO2;HO2; (g) H2S; S2−; (h) H6N22+,H6N22+, H4N2

9.

The labels are Brønsted-Lowry acid = BA; its conjugate base = CB; Brønsted-Lowry base = BB; its conjugate acid = CA. (a) HNO3(BA), H2O(BB), H3O+(CA), NO3(CB);NO3(CB); (b) CN(BB), H2O(BA), HCN(CA), OH(CB); (c) H2SO4(BA), Cl(BB), HCl(CA), HSO4(CB);HSO4(CB); (d) HSO4(BA),HSO4(BA), OH(BB), SO42−SO42−(CB), H2O(CA); (e) O2−(BB), H2O(BA) OH(CB and CA); (f) [Cu(H2O)3(OH)]+(BB), [Al(H2O)6]3+(BA), [Cu(H2O)4]2+(CA), [Al(H2O)5(OH)]2+(CB); (g) H2S(BA), NH2(BB),NH2(BB), HS(CB), NH3(CA)

11.

Amphiprotic species may either gain or lose a proton in a chemical reaction, thus acting as a base or an acid. An example is H2O. As an acid:
H2O(aq)+NH3(aq)NH4+(aq)+OH(aq).H2O(aq)+NH3(aq)NH4+(aq)+OH(aq). As a base: H2O(aq)+HCl(aq)H3O+(aq)+Cl(aq)H2O(aq)+HCl(aq)H3O+(aq)+Cl(aq)

13.

amphiprotic: (a) NH3+H3O+NH4OH+H2O,NH3+H3O+NH4OH+H2O, NH3+OCH3NH2+CH3OH;NH3+OCH3NH2+CH3OH; (b) HPO42−+OHPO43−+H2O,HPO42−+OHPO43−+H2O, HPO42−+HClO4H2PO4+ClO4;HPO42−+HClO4H2PO4+ClO4; not amphiprotic: (c) Br; (d) NH4+;NH4+; (e) AsO43−AsO43−

15.

In a neutral solution [H3O+] = [OH]. At 40 °C,
[H3O+] = [OH] = (2.910−14)1/2 = 1.7 ×× 10−7.

17.

x = 3.051 ×× 10−7 M = [H3O+] = [OH]
pH = −log3.051 ×× 10−7 = −(−6.5156) = 6.5156
pOH = pH = 6.5156

19.

(a) pH = 3.587; pOH = 10.413; (b) pH = 0.68; pOH = 13.32; (c) pOH = 3.85; pH = 10.15; (d) pH = −0.40; pOH = 14.4

21.

[H3O+] = 3.0 ×× 10−7 M; [OH] = 3.3 ×× 10−8 M

23.

[H3O+] = 1 ×× 10−2 M; [OH] = 1 ×× 10−12 M

25.

[OH] = 3.1 ×× 10−12 M

27.

The salt ionizes in solution, but the anion slightly reacts with water to form the weak acid. This reaction also forms OH, which causes the solution to be basic.

29.

[H2O] > [CH3CO2H] > [H3O+][H3O+][CH3CO2][CH3CO2] > [OH]

31.

The oxidation state of the sulfur in H2SO4 is greater than the oxidation state of the sulfur in H2SO3.

33.

Mg ( OH ) 2 ( s ) + 2HCl ( a q ) Mg 2+ ( a q ) + 2 Cl ( a q ) + 2H 2 O ( l ) BB BA CB CA Mg ( OH ) 2 ( s ) + 2HCl ( a q ) Mg 2+ ( a q ) + 2 Cl ( a q ) + 2H 2 O ( l ) BB BA CB CA

35.

Ka=2.3×10−11Ka=2.3×10−11

37.

The stronger base or stronger acid is the one with the larger Kb or Ka, respectively. In these two examples, they are (CH3)2NH and CH3NH3+.CH3NH3+.

39.

triethylamine.

41.

(a) HSO4;HSO4; higher electronegativity of the central ion. (b) H2O; NH3 is a base and water is neutral, or decide on the basis of Ka values. (c) HI; PH3 is weaker than HCl; HCl is weaker than HI. Thus, PH3 is weaker than HI. (d) PH3; in binary compounds of hydrogen with nonmetals, the acidity increases for the element lower in a group. (e) HBr; in a period, the acidity increases from left to right; in a group, it increases from top to bottom. Br is to the left and below S, so HBr is the stronger acid.

43.

(a) NaHSeO3 < NaHSO3 < NaHSO4; in polyoxy acids, the more electronegative central element—S, in this case—forms the stronger acid. The larger number of oxygen atoms on the central atom (giving it a higher oxidation state) also creates a greater release of hydrogen atoms, resulting in a stronger acid. As a salt, the acidity increases in the same manner. (b) ClO2<BrO2<IO2;ClO2<BrO2<IO2; the basicity of the anions in a series of acids will be the opposite of the acidity in their oxyacids. The acidity increases as the electronegativity of the central atom increases. Cl is more electronegative than Br, and I is the least electronegative of the three. (c) HOI < HOBr < HOCl; in a series of the same form of oxyacids, the acidity increases as the electronegativity of the central atom increases. Cl is more electronegative than Br, and I is the least electronegative of the three. (d) HOCl < HOClO < HOClO2 < HOClO3; in a series of oxyacids of the same central element, the acidity increases as the number of oxygen atoms increases (or as the oxidation state of the central atom increases). (e) HTe<HS<<PH2<NH2;HTe<HS<<PH2<NH2; PH2PH2 and NH2NH2 are anions of weak bases, so they act as strong bases toward H+. HTeHTe and HS are anions of weak acids, so they have less basic character. In a periodic group, the more electronegative element has the more basic anion. (f) BrO4<BrO3<BrO2<BrO;BrO4<BrO3<BrO2<BrO; with a larger number of oxygen atoms (that is, as the oxidation state of the central ion increases), the corresponding acid becomes more acidic and the anion consequently less basic.

45.

[ H 2 O ] > [ C 6 H 4 OH ( CO 2 H ) ] > [H + ] 0 > [C 6 H 4 OH ( CO 2 ) ] [ C 6 H 4 O ( CO 2 H ) ] > [ OH ] [ H 2 O ] > [ C 6 H 4 OH ( CO 2 H ) ] > [H + ] 0 > [C 6 H 4 OH ( CO 2 ) ] [ C 6 H 4 O ( CO 2 H ) ] > [ OH ]

47.

Strong electrolytes are 100% ionized, and, as long as the component ions are neither weak acids nor weak bases, the ionic species present result from the dissociation of the strong electrolyte. Equilibrium calculations are necessary when one (or more) of the ions is a weak acid or a weak base.

49.

1. Assume that the change in initial concentration of the acid as the equilibrium is established can be neglected, so this concentration can be assumed constant and equal to the initial value of the total acid concentration. 2. Assume we can neglect the contribution of water to the equilibrium concentration of H3O+.

51.

(b) The addition of HCl

53.

(a) Adding HCl will add H3O+ ions, which will then react with the OH ions, lowering their concentration. The equilibrium will shift to the right, increasing the concentration of HNO2, and decreasing the concentration of NO2NO2 ions. (b) Adding HNO2 increases the concentration of HNO2 and shifts the equilibrium to the left, increasing the concentration of NO2NO2 ions and decreasing the concentration of OH ions. (c) Adding NaOH adds OH ions, which shifts the equilibrium to the left, increasing the concentration of NO2NO2 ions and decreasing the concentrations of HNO2. (d) Adding NaCl has no effect on the concentrations of the ions. (e) Adding KNO2 adds NO2NO2 ions and shifts the equilibrium to the right, increasing the HNO2 and OH ion concentrations.

55.

This is a case in which the solution contains a mixture of acids of different ionization strengths. In solution, the HCO2H exists primarily as HCO2H molecules because the ionization of the weak acid is suppressed by the strong acid. Therefore, the HCO2H contributes a negligible amount of hydronium ions to the solution. The stronger acid, HCl, is the dominant producer of hydronium ions because it is completely ionized. In such a solution, the stronger acid determines the concentration of hydronium ions, and the ionization of the weaker acid is fixed by the [H3O+] produced by the stronger acid.

57.

(a) Kb=1.8×10−5;Kb=1.8×10−5;
(b) Ka=4.5×10−4;Ka=4.5×10−4;
(c) Kb=7.4×10−5;Kb=7.4×10−5;
(d) Ka=5.6×10−10Ka=5.6×10−10

59.

K a = 1.2 × 10 −2 K a = 1.2 × 10 −2

61.

(a) Kb=4.3×10−12;Kb=4.3×10−12;
(b) Ka=1.6×10−8;Ka=1.6×10−8;
(c) Kb=5.9×10−7;Kb=5.9×10−7;
(d) Kb=4.2×10−3;Kb=4.2×10−3;
(e) Kb=2.3×10−3;Kb=2.3×10−3;
(f) Kb=6.3×10−13Kb=6.3×10−13

63.

(a) is the correct statement.

65.

[H3O+] = 7.5 ×× 10−3 M
[HNO2] = 0.127
[OH] = 1.3 ×× 10−12 M
[BrO] = 4.5 ×× 10−8 M
[HBrO] = 0.120 M

67.

[OH] = [NO4+][NO4+] = 0.0014 M
[NH3] = 0.144 M
[H3O+] = 6.9 ×× 10−12 M
[C6H5NH3+][C6H5NH3+] = 3.9 ×× 10−8 M
[C6H5NH2] = 0.100 M

69.

(a) [H3O+][ClO][HClO]=(x)(x)(0.0092x)(x)(x)0.0092=2.9×10−8[H3O+][ClO][HClO]=(x)(x)(0.0092x)(x)(x)0.0092=2.9×10−8
Solving for x gives 1.63 ×× 10−5 M. This value is less than 5% of 0.0092, so the assumption that it can be neglected is valid. Thus, the concentrations of solute species at equilibrium are:
[H3O+] = [ClO] = 5.8 ×× 10−5 M
[HClO] = 0.00092 M
[OH] = 6.1 ×× 10−10 M;
(b) [C6H5NH3+][OH][C6H5NH2]=(x)(x)(0.0784x)(x)(x)0.0784=4.3×10−10[C6H5NH3+][OH][C6H5NH2]=(x)(x)(0.0784x)(x)(x)0.0784=4.3×10−10
Solving for x gives 5.81 ×× 10−6 M. This value is less than 5% of 0.0784, so the assumption that it can be neglected is valid. Thus, the concentrations of solute species at equilibrium are:
[CH3CO2][CH3CO2] = [OH] = 5.8 ×× 10−6 M
[C6H5NH2] = 0.00784
[H3O+] = 1.7×× 10−9 M;
(c) [H3O+][CN][HCN]=(x)(x)(0.0810x)(x)(x)0.0810=4.9×10−10[H3O+][CN][HCN]=(x)(x)(0.0810x)(x)(x)0.0810=4.9×10−10
Solving for x gives 6.30 ×× 10−6 M. This value is less than 5% of 0.0810, so the assumption that it can be neglected is valid. Thus, the concentrations of solute species at equilibrium are:
[H3O+] = [CN] = 6.3 ×× 10−6 M
[HCN] = 0.0810 M
[OH] = 1.6 ×× 10−9 M;
(d) [(CH3)3NH+][OH][(CH3)3N]=(x)(x)(0.11x)(x)(x)0.11=6.3×10−5[(CH3)3NH+][OH][(CH3)3N]=(x)(x)(0.11x)(x)(x)0.11=6.3×10−5
Solving for x gives 2.63 ×× 10−3 M. This value is less than 5% of 0.11, so the assumption that it can be neglected is valid. Thus, the concentrations of solute species at equilibrium are:
[(CH3)3NH+] = [OH] = 2.6 ×× 10−3 M
[(CH3)3N] = 0.11 M
[H3O+] = 3.8 ×× 10−12 M;
(e) [Fe(H2O)5(OH)+][H3O+][Fe(H2O)62+]=(x)(x)(0.120x)(x)(x)0.120=1.6×10−7[Fe(H2O)5(OH)+][H3O+][Fe(H2O)62+]=(x)(x)(0.120x)(x)(x)0.120=1.6×10−7
Solving for x gives 1.39 ×× 10−4 M. This value is less than 5% of 0.120, so the assumption that it can be neglected is valid. Thus, the concentrations of solute species at equilibrium are:
[Fe(H2O)5(OH)+] = [H3O+] = 1.4 ×× 10−4 M
[Fe(H2O)62+][Fe(H2O)62+] = 0.120 M
[OH] = 7.2 ×× 10−11 M

71.

pH = 2.41

73.

[C10H14N2] = 0.049 M
[C10H14N2H+] = 1.9 ×× 10−4 M
[C10H14N2H22+][C10H14N2H22+] = 1.4 ×× 10−11 M
[OH] = 1.9 ×× 10−4 M
[H3O+] = 5.3 ×× 10−11 M

75.

K a = 1.2 × 10 −2 K a = 1.2 × 10 −2

77.

K b = 1.77 × 10 −5 K b = 1.77 × 10 −5

79.

(a) acidic; (b) basic; (c) acidic; (d) neutral

81.

[H3O+] and [HCO3][HCO3] are practically equal

83.

[C6H4(CO2H)2] 7.2 ×× 10−3 M, [C6H4(CO2H)(CO2)] = [H3O+] 2.8 ×× 10−3 M, [C6H4(CO2)22−][C6H4(CO2)22−]3.9 ×× 10−6 M, [OH] 3.6 ×× 10−12 M

85.

(a) Ka2=1.5×10−11;Ka2=1.5×10−11;
(b) Kb=4.3×10−12;Kb=4.3×10−12;
(c) [Te2−][H3O+][HTe]=(x)(0.0141+x)(0.0141x)(x)(0.0141)0.0141=1.5×10−11[Te2−][H3O+][HTe]=(x)(0.0141+x)(0.0141x)(x)(0.0141)0.0141=1.5×10−11
Solving for x gives 1.5 ×× 10−11 M. Therefore, compared with 0.014 M, this value is negligible (1.1 ×× 10−7%).

87.

Excess H3O+ is removed primarily by the reaction:
H3O+(aq)+H2PO4(aq)H3PO4(aq)+H2O(l)H3O+(aq)+H2PO4(aq)H3PO4(aq)+H2O(l)
Excess base is removed by the reaction:
OH(aq)+H3PO4(aq)H2PO4(aq)+H2O(l)OH(aq)+H3PO4(aq)H2PO4(aq)+H2O(l)

89.

[H3O+] = 1.5 ×× 10−4 M

91.

[OH] = 4.2 ×× 10−4 M

93.

[NH4NO3] = 0.36 M

95.

(a) The added HCl will increase the concentration of H3O+ slightly, which will react with CH3CO2CH3CO2 and produce CH3CO2H in the process. Thus, [CH3CO2][CH3CO2] decreases and [CH3CO2H] increases.
(b) The added KCH3CO2 will increase the concentration of [CH3CO2][CH3CO2] which will react with H3O+ and produce CH3CO2 H in the process. Thus, [H3O+] decreases slightly and [CH3CO2H] increases.
(c) The added NaCl will have no effect on the concentration of the ions.
(d) The added KOH will produce OH ions, which will react with the H3O+, thus reducing [H3O+]. Some additional CH3CO2H will dissociate, producing [CH3CO2][CH3CO2] ions in the process. Thus, [CH3CO2H] decreases slightly and [CH3CO2][CH3CO2] increases.
(e) The added CH3CO2H will increase its concentration, causing more of it to dissociate and producing more [CH3CO2][CH3CO2] and H3O+ in the process. Thus, [H3O+] increases slightly and [CH3CO2][CH3CO2] increases.

97.

pH = 8.95

99.

37 g (0.27 mol)

101.

(a) pH = 5.222;
(b) The solution is acidic.
(c) pH = 5.221

103.

To prepare the best buffer for a weak acid HA and its salt, the ratio [H3O+]Ka[H3O+]Ka should be as close to 1 as possible for effective buffer action. The [H3O+] concentration in a buffer of pH 3.1 is [H3O+] = 10−3.1 = 7.94 ×× 10−4 M
We can now solve for Ka of the best acid as follows:
[H3O+]Ka=1Ka=[H3O+]1=7.94×10−4[H3O+]Ka=1Ka=[H3O+]1=7.94×10−4
In Table 14.2, the acid with the closest Ka to 7.94 ×× 10−4 is HF, with a Ka of 7.2 ×× 10−4.

105.

For buffers with pHs > 7, you should use a weak base and its salt. The most effective buffer will have a ratio [OH]Kb[OH]Kb that is as close to 1 as possible. The pOH of the buffer is 14.00 − 10.65 = 3.35. Therefore, [OH] is [OH] = 10−pOH = 10−3.35 = 4.467 ×× 10−4 M.
We can now solve for Kb of the best base as follows:
[OH]Kb=1[OH]Kb=1
Kb = [OH] = 4.47 ×× 10−4
In Table 14.3, the base with the closest Kb to 4.47 ×× 10−4 is CH3NH2, with a Kb = 4.4 ×× 10−4.

107.

The molar mass of sodium saccharide is 205.169 g/mol. Using the abbreviations HA for saccharin and NaA for sodium saccharide the number of moles of NaA in the solution is:
9.75 ×× 10−6 mol
The pKa for [HA] is 1.68, so [HA] = 6.2×19–96.2×19–9 M. Thus, [A–] (saccharin ions) is 3.90×10–53.90×10–5 M.
Thus, [A](saccarin ions) is 3.90 ×× 10−5 M

109.

At the equivalence point in the titration of a weak base with a strong acid, the resulting solution is slightly acidic due to the presence of the conjugate acid. Thus, pick an indicator that changes color in the acidic range and brackets the pH at the equivalence point. Methyl orange is a good example.

111.

In an acid solution, the only source of OH ions is water. We use Kw to calculate the concentration. If the contribution from water was neglected, the concentration of OH would be zero.

113.


A graph is shown that is titled “Plot of [ H subscript 3 O superscript + ] Against [ H F ].” The horizontal axis is labeled “[ H F ], M.” The axis begins at 10 superscript negative 10 and includes markings every 10 superscript 2 units up to 1.0. The vertical axis is labeled “[ H subscript 3 O superscript plus ], M” and begins at 10 superscript negative 10 and increases by 10 superscript 2 up to 1.0. A black curve starts at the left side of the graph at (10 superscript negative 10, 10 superscript negative 7). The line extends horizontally to a horizontal axis value of 10 superscript negative 8. After this, the line gradually increases at a steady rate to a value just over 10 superscript negative 3 at a horizontal axis value of 10 superscript negative 2.
115.

(a) pH = 2.50;
(b) pH = 4.01;
(c) pH = 5.60;
(d) pH = 8.35;
(e) pH = 11.08

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