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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. 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 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 Electrochemistry
    1. Introduction
    2. 16.1 Review of Redox Chemistry
    3. 16.2 Galvanic Cells
    4. 16.3 Electrode and Cell Potentials
    5. 16.4 Potential, Free Energy, and Equilibrium
    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.

(a) AgI(s)Ag+(aq)+I(aq)xx_AgI(s)Ag+(aq)+I(aq)xx_
(b) CaCO3(s)Ca2+(aq)+CO32−(aq)x_xCaCO3(s)Ca2+(aq)+CO32−(aq)x_x
(c) Mg(OH)2(s)Mg2+(aq)+2OH(aq)x2x_Mg(OH)2(s)Mg2+(aq)+2OH(aq)x2x_
(d) Mg3(PO4)2(s)3Mg2+(aq)+2PO43−(aq)3x_2xMg3(PO4)2(s)3Mg2+(aq)+2PO43−(aq)3x_2x
(e) Ca5(PO4)3OH(s)5Ca2+(aq)+3PO43−(aq)+OH(aq)5x_3x_xCa5(PO4)3OH(s)5Ca2+(aq)+3PO43−(aq)+OH(aq)5x_3x_x

3.

There is no change. A solid has an activity of 1 whether there is a little or a lot.

5.

The solubility of silver bromide at the new temperature must be known. Normally the solubility increases and some of the solid silver bromide will dissolve.

7.

CaF2, MnCO3, and ZnS

9.

(a) LaF3(s)La3+(aq)+3F(aq)Ksp=[La3+][F]3;LaF3(s)La3+(aq)+3F(aq)Ksp=[La3+][F]3;
(b) CaCO3(s)Ca2+(aq)+CO32−(aq)Ksp=[Ca2+][CO32−];CaCO3(s)Ca2+(aq)+CO32−(aq)Ksp=[Ca2+][CO32−];
(c) Ag2SO4(s)2Ag+(aq)+SO42−(aq)Ksp=[Ag+]2[SO42−];Ag2SO4(s)2Ag+(aq)+SO42−(aq)Ksp=[Ag+]2[SO42−];
(d) Pb(OH)2(s)Pb2+(aq)+2OH(aq)Ksp=[Pb2+][OH]2Pb(OH)2(s)Pb2+(aq)+2OH(aq)Ksp=[Pb2+][OH]2

11.

(a)1.77 ×× 10–7; (b) 1.6 ×× 10–6; (c) 2.2 ×× 10–9; (d) 7.91 ×× 10–22

13.

(a) 2 ×× 10–2 M; (b) 1.5 ×× 10–3 M; (c) 2.27 ×× 10–9 M; (d) 2.2 ×× 10–10 M

15.

(a) 6.4 ×× 10−9 M = [Ag+], [Cl] = 0.025 M. Check: 6.4×109M0.025M×100%=2.6×105%,6.4×109M0.025M×100%=2.6×105%,an insignificant change;
(b) 2.2 ×× 10−5 M = [Ca2+], [F] = 0.0013 M. Check: 2.26×105M0.00133M×100%=1.70%.2.26×105M0.00133M×100%=1.70%. This value is less than 5% and can be ignored.
(c) 0.2238 M = [SO42];[SO42]; [Ag+] = 7.4 ×× 10–3 M. Check: 3.7×1030.2238×100%=1.64×102;3.7×1030.2238×100%=1.64×102; the condition is satisfied.
(d) [OH] = 2.8 ×× 10–3 M; 5.7 ×× 10−12 M = [Zn2+]. Check: 5.7×10122.8×103×100%=2.0×107%;5.7×10122.8×103×100%=2.0×107%; x is less than 5% of [OH] and is, therefore, negligible.

17.

(a) [Cl] = 7.6 ×× 10−3 M
Check: 7.6×1030.025×100%=30%7.6×1030.025×100%=30%
This value is too large to drop x. Therefore solve by using the quadratic equation:
[Ti+] = 3.1 ×× 10–2 M
[Cl] = 6.1 ×× 10–3
(b) [Ba2+] = 7.7 ×× 10–4 M
Check: 7.7×1040.0313×100%=2.4%7.7×1040.0313×100%=2.4%
Therefore, the condition is satisfied.
[Ba2+] = 7.7 ×× 10–4 M
[F] = 0.0321 M;
(c) Mg(NO3)2 = 0.02444 M
[C2O42−]=2.9×10−5[C2O42−]=2.9×10−5
Check: 2.9×10−50.02444×100%=0.12%2.9×10−50.02444×100%=0.12%
The condition is satisfied; the above value is less than 5%.
[C2O42−]=2.9×10−5M[C2O42−]=2.9×10−5M
[Mg2+] = 0.0244 M
(d) [OH] = 0.0501 M
[Ca2+] = 3.15 ×× 10–3
Check: 3.15×10−30.050×100%=6.28%3.15×10−30.050×100%=6.28%
This value is greater than 5%, so a more exact method, such as successive approximations, must be used.
[Ca2+] = 2.8 ×× 10–3 M
[OH] = 0.053 ×× 10–2 M

19.

The changes in concentration are greater than 5% and thus exceed the maximum value for disregarding the change.

21.

CaSO4∙2H2O is the most soluble Ca salt in mol/L, and it is also the most soluble Ca salt in g/L.

23.

4.8 ×× 10–3 M = [SO42−][SO42−] = [Ca2+]; Since this concentration is higher than 2.60 ×× 10–3 M, “gyp” water does not meet the standards.

25.

Mass (CaSO4·2H2O) = 0.72 g/L

27.

(a) [Ag+] = [I] = 1.3 ×× 10–5 M; (b) [Ag+] = 2.88 ×× 10–2 M, [SO42−][SO42−]= 1.44 ×× 10–2 M; (c) [Mn2+] = 3.7 ×× 10–5 M, [OH] = 7.4 ×× 10–5 M; (d) [Sr2+] = 4.3 ×× 10–2 M, [OH] = 8.6 ×× 10–2 M; (e) [Mg2+] = 1.3 ×× 10–4 M, [OH] = 2.6 ×× 10–4 M.

29.

(a) 1.45 ×× 10–4; (b) 8.2 ×× 10–55; (c) 1.35 ×× 10–4; (d) 1.18 ×× 10–5; (e) 1.08 ×× 10–10

31.

(a) CaCO3 does precipitate. (b) The compound does not precipitate. (c) The compound does not precipitate. (d) The compound precipitates.

33.

3.03 ×× 10−7 M

35.

9.2 ×× 10−13 M

37.

[Ag+] = 1.8 ×× 10–3 M

39.

6.3 ×× 10–4

41.

(a) 2.25 L; (b) 7.2 ×× 10–7 g

43.

100% of it is dissolved

45.

(a) Hg22+Hg22+ and Cu2+: Add SO42−.SO42−. (b) SO42−SO42− and Cl: Add Ba2+. (c) Hg2+ and Co2+: Add S2–. (d) Zn2+ and Sr2+: Add OH until [OH] = 0.050 M. (e) Ba2+ and Mg2+: Add SO42−.SO42−. (f) CO32−CO32− and OH: Add Ba2+.

47.

AgI will precipitate first.

49.

1.5 ×× 10−12 M

51.

3.99 kg

53.

(a) 3.1 ×× 10–11; (b) [Cu2+] = 2.6 ×× 10–3; [IO3][IO3] = 5.3 ×× 10–3

55.

1.8 ×× 10–5 g Pb(OH)2

57.

Mg(OH)2(s)Mg2++2OHKsp=[Mg2+][OH]2Mg(OH)2(s)Mg2++2OHKsp=[Mg2+][OH]2
1.23 ×× 10−3 g Mg(OH)2

59.

MnCO3 will form first, since it has the smallest Ksp value it is the least soluble. MnCO3 will be the last to precipitate, it has the largest Ksp value.

62.

when the amount of solid is so small that a saturated solution is not produced

64.

8 ×× 10–5 M

66.

5 ×× 1023

68.
This table has two main columns and three rows. The first row for the first column does not have a heading and then has the following in the first column: Initial concentration ( M ) and Equilibrium ( M ). The second column has the header, “[ C d ( C N ) subscript 4 to the second power superscript negative sign ] [ C N superscript negative sign ] [ C d to the second power superscript positive sign ].” Under the second column is a subgroup of two rows and three columns. The first column contains the following: 0.250 and 0.250 minus x. The second column contains the following: 0 and 4 x. The third column contains the following: 0 and x.


[Cd2+] = 9.5 ×× 10–5 M; [CN] = 3.8 ×× 10–4 M

70.

[Co3+] = 3.0 ×× 10–6 M; [NH3] = 1.8 ×× 10–5 M

72.

1.3 g

74.

0.79 g

76.

(a)

This figure shows a chemical reaction modeled with structural formulas. On the left side is a structure with a central C atom. O atoms, each with two unshared electron pairs, are double bonded to the left and right sides of the C atom. Following a plus sign is another structure in brackets which has an O atom with three unshared electron dot pairs single bonded to an H atom on the right. Outside the brackets is superscript negative sign. Following a right pointing arrow is a structure in brackets that has a central C atom to which 3 O atoms are bonded. Above and slightly to the right, one of the O atoms is connected with a double bond. This O atom has two unshared electron pairs. The second O atom is single bonded below and slightly to the right. This O atom has three unshared electron pairs. The third O atom is bonded to the left of the C atom. This O atom has two unshared electron pairs and an H atom single bonded to its left. Outside the brackets to the right is a superscript negative symbol.


(b)

This figure shows a chemical reaction modeled with structural formulas. On the left side is a structure that has a central B atom to which 3 O atoms are bonded. The O atoms above and below slightly right of the B atom each have an H atom single bonded to the right. The third O atom is single bonded to the left side of the B atom. This O atom has an H atom single bonded to its left side. All O atoms in this structure have two unshared electron pairs. Following a plus sign is another structure which has an O atom single bonded to an H atom on its right. The O atom has three unshared electron pairs. The structure appears in brackets with a superscript negative sign. Following a right pointing arrow is a structure in brackets has a central B atom to which 4 O atoms are bonded. The O atoms above, below, and right of the B atom each hav an H atom single bonded to the right. The third O atom is single bonded to the left side of the B atom. This O atom has an H atom single bonded to its left side. All O atoms in this structure have two unshared electron pairs. Outside the brackets to the right is a superscript negative symbol.


(c)

This figure illustrates a chemical reaction using structural formulas. On the left, two I atoms, each with 3 unshared electron pairs, are joined with a single bond. Following a plus sign is another structure which has an I atom with four pairs of electron dots and a superscript negative sign. Following a right pointing arrow is a structure in brackets that has three I atoms connected in a line with single bonds. The two end I atoms have three unshared electron dot pairs and the I atom at the center has two unshared electron pairs. Outside the brackets is a superscript negative sign.


(d)

This figure illustrates a chemical reaction using structural formulas. On the left, an A l atom is positioned at the center of a structure and three Cl atoms are single bonded above, left, and below. Each C l atom has three pairs of electron dots. Following a plus sign is another structure which has an F atom is surrounded by four electron dot pairs and a superscript negative symbol. Following a right pointing arrow is a structure in brackets that has a central A l atom to which 4 C l atoms are connected with single bonds above, below, to the left, and to the right. Each C l atom in this structure has three pairs of electron dots. Outside the brackets is a superscript negative symbol.


(e)

This figure illustrates a chemical reaction using structural formulas. On the left is a structure which has an S atom at the center. O atoms are single bonded above and below. These O atoms have three electron dot pairs each. To the right of the S atom is a double bonded O atom which has two pairs of electron dots. Following a plus sign is an O atom which is surrounded by four electron dot pairs and has a superscript 2 negative. Following a right pointing arrow is a structure in brackets that has a central S atom to which 4 O atoms are connected with single bonds above, below, to the left, and to the right. Each of the O atoms has three pairs of electron dots. Outside the brackets is a superscript 2 negative.
78.

(a)

This figure represents a chemical reaction in two rows. The top row shows the reaction using chemical formulas. The second row uses structural formulas to represent the reaction. The first row contains the equation H C l ( g ) plus P H subscript 3 ( g ) right pointing arrow left bracket P H subscript 4 right bracket superscript plus plus left bracket C l with 4 pairs of electron dots right bracket superscript negative sign. The second row begins on the left with H left bracket C l with four unshared electron pairs right bracket plus a structure in brackets with a central P atom with H atoms single bonded at the left, above, and to the right. A single unshared electron pair is on the central P atom. Outside the brackets to the right is a superscript plus sign. Following a right pointing arrow is a structure in brackets with a central P atom with H atoms single bonded at the left, above, below, and to the right. Outside the brackets is a superscript plus sign. This structure is followed by a plus and a C l atom in brackets with four unshared electron pairs and a superscript negative sign.


(b) H3O++CH3CH4+H2OH3O++CH3CH4+H2O

This figure represents a chemical reaction using structural formulas. A structure is shown in brackets on the left which is composed of a central O atom with one unshared electron pair and three single bonded H atoms to the left, right, and above the atom. Outside the brackets to the right is a superscript plus sign. Following a plus sign, is another structure in brackets composed of a central C atom with one unshared electron pair and three single bonded H atoms to the left, right, and above the atom. Outside the brackets to the right is a superscript negative sign. Following a right pointing arrow is a structure with a central C atom with H atoms single bonded above, below, left and right. Following a plus sign is a structure with a central O atom with two unshared electron pairs and two H atoms connected with single bonds.


(c) CaO+SO3CaSO4CaO+SO3CaSO4

This figure represents a chemical reaction using structural formulas. On the left, C a superscript 2 plus is just left of bracket O with four unshared electron pairs right bracket superscript 2 negative plus a structure with a central S atom to which two O atoms are single bonded at the left and right, and a single O atom is double bonded above. The two single bonded O atoms each have three unshared electron pairs and the double bonded O atom has two unshared electron pairs. Following a right pointing arrow is C a superscript 2 plus just left of a structure in brackets with a central S atom which has 4 O atoms single bonded at the left, above, below, and to the right. Each of the O atoms has three unshared electron pairs. Outside the brackets to the right is a superscript two negative.


(d) NH4++C2H5OC2H5OH+NH3NH4++C2H5OC2H5OH+NH3

This figure represents a chemical reaction using structural formulas. A structure is shown in brackets on the left which is composed of a central N atom with four single bonded H atoms to the left, right, above, and below the atom. Outside the brackets to the right is a superscript plus sign. Following a plus sign, is another structure in brackets composed of a C atom with three single bonded H atoms above, below, and to the left. A second C atom is single bonded to the right. This C atom has H atoms single bonded above and below. To the right of the second C atom, an O atom is single bonded. This O atom has three unshared electron pairs. Outside the brackets to the right is a subperscript negative. Following a right pointing arrow is a structure composed of a C atom with three single bonded H atoms above, below, and to the left. A second C atom is single bonded to the right. This C atom has H atoms single bonded above and below. To the right of the second C atom, an O atom is single bonded. This O atom has two unshared electron pairs and an H atom single bonded to its right.
80.

0.0281 g

82.

HNO3(l)+HF(l)H2NO3++F;HNO3(l)+HF(l)H2NO3++F; HF(l)+BF3(g)H++BF4HF(l)+BF3(g)H++BF4

84.

(a) H3BO3+H2OH4BO4+H+;H3BO3+H2OH4BO4+H+; (b) The electronic and molecular shapes are the same—both tetrahedral. (c) The tetrahedral structure is consistent with sp3 hybridization.

86.

0.014 M

88.

7.2 ×× 10–15 M

90.

4.4 ×× 10−22 M

93.

[OH] = 4.5 ×× 10−5; [Al3+] = 2.2 ×× 10–20 (molar solubility)

95.

[SO42−]=0.049M[SO42−]=0.049M; [Ba2+] = 4.7 ×× 10–7 (molar solubility)

97.

[OH] = 7.6 ×× 10−3 M; [Pb2+] = 2.1 ×× 10–11 (molar solubility)

99.

7.66

101.

(a) Ksp = [Mg2+][F]2 = (1.21 ×× 10–3)(2 ×× 1.21 ×× 10–3)2 = 7.09 ×× 10–9
(b) 7.09 ×× 10–7 M
(c) Determine the concentration of Mg2+ and F that will be present in the final volume. Compare the value of the ion product [Mg2+][F]2 with Ksp. If this value is larger than Ksp, precipitation will occur.
0.1000 L ×× 3.00 ×× 10–3 M Mg(NO3)2 = 0.3000 L ×× M Mg(NO3)2
M Mg(NO3)2 = 1.00 ×× 10–3 M
0.2000 L ×× 2.00 ×× 10–3 M NaF = 0.3000 L ×× M NaF
M NaF = 1.33 ×× 10–3 M
ion product = (1.00 ×× 10–3)(1.33 ×× 10–3)2 = 1.77 ×× 10–9 This value is smaller than Ksp, so no precipitation will occur.
(d) MgF2 is less soluble at 27 °C than at 18 °C. Because added heat acts like an added reagent, when it appears on the product side, the Le Châtelier’s principle states that the equilibrium will shift to the reactants’ side to counter the stress. Consequently, less reagent will dissolve. This situation is found in our case. Therefore, the reaction is exothermic.

103.

BaF2, Ca3(PO4)2, ZnS; each is a salt of a weak acid, and the [H3O+][H3O+] from perchloric acid reduces the equilibrium concentration of the anion, thereby increasing the concentration of the cations

105.

Effect on amount of solid CaHPO4, [Ca2+], [OH]: (a) increase, increase, decrease; (b) decrease, increase, decrease; (c) no effect, no effect, no effect; (d) decrease, increase, decrease; (e) increase, no effect, no effect

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© Feb 14, 2019 OpenStax. Textbook content produced by OpenStax is licensed under a Creative Commons Attribution License 4.0 license. The OpenStax name, OpenStax logo, OpenStax book covers, OpenStax CNX name, and OpenStax CNX logo are not subject to the Creative Commons license and may not be reproduced without the prior and express written consent of Rice University.