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Chemistry 2e

Chapter 14

Chemistry 2eChapter 14
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  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 Molecular and Ionic 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
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 × 10−14)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(aq)Mg2+(aq)+2Cl(aq)+2H2O(l)BBBACBCAMg(OH)2(s)+2HCl(aq)Mg2+(aq)+2Cl(aq)+2H2O(l)BBBACBCA

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.

[H2O]>[C6H4OH(CO2H)]>[H+]0>[C6H4OH(CO2)][C6H4O(CO2H)]>[OH][H2O]>[C6H4OH(CO2H)]>[H+]0>[C6H4OH(CO2)][C6H4O(CO2H)]>[OH]

47.

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+.

48.

(b) The addition of HCl

50.

(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.

52.

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.

54.

(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

56.

Ka=1.2×10−2Ka=1.2×10−2

58.

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

60.

(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

62.

pH = 2.41

64.

[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

66.

Ka=1.2×10−2Ka=1.2×10−2

68.

Kb=1.77×10−5Kb=1.77×10−5

70.

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

72.

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

74.

[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

76.

(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%).

78.

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)

80.

[H3O+] = 1.5 ×× 10−4 M

82.

[OH] = 4.2 ×× 10−4 M

84.

(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.

86.

pH = 8.95

88.

37 g (0.27 mol)

90.

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

92.

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.

94.

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

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