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Chemistry 2e

Chapter 18

Chemistry 2eChapter 18
  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.

The alkali metals all have a single s electron in their outermost shell. In contrast, the alkaline earth metals have a completed s subshell in their outermost shell. In general, the alkali metals react faster and are more reactive than the corresponding alkaline earth metals in the same period.

3.


Na+I22NaI2Na+SeNa2Se2Na+O2Na2O2Na+I22NaI2Na+SeNa2Se2Na+O2Na2O2
Sr+I2SrI2Sr+SeSrSe2Sr+O22SrOSr+I2SrI2Sr+SeSrSe2Sr+O22SrO
2Al+3I22AlI32Al+3SeAl2Se34Al+3O22Al2O32Al+3I22AlI32Al+3SeAl2Se34Al+3O22Al2O3

5.

The possible ways of distinguishing between the two include infrared spectroscopy by comparison of known compounds, a flame test that gives the characteristic yellow color for sodium (strontium has a red flame), or comparison of their solubilities in water. At 20 °C, NaCl dissolves to the extent of 35.7 g100 mL35.7 g100 mL compared with 53.8 g100 mL53.8 g100 mL for SrCl2. Heating to 100 °C provides an easy test, since the solubility of NaCl is 39.12 g100 mL,39.12 g100 mL, but that of SrCl2 is 100.8 g100 mL.100.8 g100 mL. Density determination on a solid is sometimes difficult, but there is enough difference (2.165 g/mL NaCl and 3.052 g/mL SrCl2) that this method would be viable and perhaps the easiest and least expensive test to perform.

7.

(a) 2Sr(s)+O2(g)2SrO(s);2Sr(s)+O2(g)2SrO(s); (b) Sr(s)+2HBr(g)SrBr2(s)+H2(g);Sr(s)+2HBr(g)SrBr2(s)+H2(g); (c) Sr(s)+H2(g)SrH2(s);Sr(s)+H2(g)SrH2(s); (d) 6Sr(s)+P4(s)2Sr3P2(s);6Sr(s)+P4(s)2Sr3P2(s); (e) Sr(s)+2H2O(l)Sr(OH)2(aq)+H2(g)Sr(s)+2H2O(l)Sr(OH)2(aq)+H2(g)

9.

11 lb

11.

Yes, tin reacts with hydrochloric acid to produce hydrogen gas.

13.

In PbCl2, the bonding is ionic, as indicated by its melting point of 501 °C. In PbCl4, the bonding is covalent, as evidenced by it being an unstable liquid at room temperature.

15.

2CsCl(l)+Ca(g)countercurrentfractionatingtower2Cs(g)+CaCl2(l)2CsCl(l)+Ca(g)countercurrentfractionatingtower2Cs(g)+CaCl2(l)

17.

Cathode (reduction): 2Li++2e2Li(l);2Li++2e2Li(l); Anode (oxidation): 2ClCl2(g)+2e;2ClCl2(g)+2e; Overall reaction: 2Li++2Cl2Li(l)+Cl2(g)2Li++2Cl2Li(l)+Cl2(g)

19.

0.5035 g H2

21.

Despite its reactivity, magnesium can be used in construction even when the magnesium is going to come in contact with a flame because a protective oxide coating is formed, preventing gross oxidation. Only if the metal is finely subdivided or present in a thin sheet will a high-intensity flame cause its rapid burning.

23.

Extract from ore: AlO(OH)(s)+NaOH(aq)+H2O(l)Na[ Al(OH)4 ](aq)AlO(OH)(s)+NaOH(aq)+H2O(l)Na[ Al(OH)4 ](aq)
Recover: 2Na[ Al(OH)4 ](s)+H2SO4(aq)2Al(OH)3(s)+Na2SO4(aq)+2H2O(l)2Na[ Al(OH)4 ](s)+H2SO4(aq)2Al(OH)3(s)+Na2SO4(aq)+2H2O(l)
Sinter: 2Al(OH)3(s)Al2O3(s)+3H2O(g)2Al(OH)3(s)Al2O3(s)+3H2O(g)
Dissolve in Na3AlF6(l) and electrolyze: Al3++3eAl(s)Al3++3eAl(s)

25.

25.83%

27.

39 kg

29.

(a) H3BPH3:

This Lewis structure is composed of a boron atom single bonded to a phosphorus atom. Each of these atoms is single bonded to three hydrogen atoms.


(b) BF4:BF4:

This Lewis structure is composed of a boron atom single bonded to four fluorine atoms, each of which has three lone pairs of electrons. The structure is surrounded by brackets, and a negative sign appears as a superscript outside the brackets.


(c) BBr3:

This Lewis structure is composed of a boron atom single bonded to three bromine atoms, each of which has three lone pairs of electrons.


(d) B(CH3)3:

This Lewis structure is composed of a boron atom that is single bonded to three carbon atoms, each of which is single bonded to three hydrogen atoms.


(e) B(OH)3:

This Lewis structure is composed of a boron atom that is single bonded to three oxygen atoms, each of which has two lone pairs of electrons. Each oxygen atom is single bonded to a hydrogen atom.
31.

1s22s22p63s23p23d0.

33.

(a) (CH3)3SiH: sp3 bonding about Si; the structure is tetrahedral; (b) SiO44−:SiO44−: sp3 bonding about Si; the structure is tetrahedral; (c) Si2H6: sp3 bonding about each Si; the structure is linear along the Si-Si bond; (d) Si(OH)4: sp3 bonding about Si; the structure is tetrahedral; (e) SiF62−:SiF62−: sp3d2 bonding about Si; the structure is octahedral

35.

(a) nonpolar; (b) nonpolar; (c) polar; (d) nonpolar; (e) polar

37.

(a) tellurium dioxide or tellurium(IV) oxide; (b) antimony(III) sulfide; (c) germanium(IV) fluoride; (d) silane or silicon(IV) hydride; (e) germanium(IV) hydride

39.

Boron has only s and p orbitals available, which can accommodate a maximum of four electron pairs. Unlike silicon, no d orbitals are available in boron.

41.

(a) ΔH° = 87 kJ; ΔG° = 44 kJ; (b) ΔH° = −109.9 kJ; Δ = −154.7 kJ; (c) ΔH° = −510 kJ; ΔG° = −601.5 kJ

43.

A mild solution of hydrofluoric acid would dissolve the silicate and would not harm the diamond.

45.

In the N2 molecule, the nitrogen atoms have an σ bond and two π bonds holding the two atoms together. The presence of three strong bonds makes N2 a very stable molecule. Phosphorus is a third-period element, and as such, does not form π bonds efficiently; therefore, it must fulfill its bonding requirement by forming three σ bonds.

47.

(a) H = 1+, C = 2+, and N = 3−; (b) O = 2+ and F = 1−; (c) As = 3+ and Cl = 1−

49.

S < Cl < O < F

51.

The electronegativity of the nonmetals is greater than that of hydrogen. Thus, the negative charge is better represented on the nonmetal, which has the greater tendency to attract electrons in the bond to itself.

53.

Hydrogen has only one orbital with which to bond to other atoms. Consequently, only one two-electron bond can form.

55.

0.43 g H2

57.

(a) Ca(OH)2(aq)+CO2(g)CaCO3(s)+H2O(l);Ca(OH)2(aq)+CO2(g)CaCO3(s)+H2O(l); (b) CaO(s)+SO2(g)CaSO3(s);CaO(s)+SO2(g)CaSO3(s);
(c) 2NaHCO3(s)+NaH2PO4(aq)Na3PO4(aq)+2CO2(g)+2H2O(l)2NaHCO3(s)+NaH2PO4(aq)Na3PO4(aq)+2CO2(g)+2H2O(l)

59.

(a) NH2−:

This Lewis structure shows a nitrogen atom with three lone pairs of electrons single bonded to a hydrogen atom. The structure is surrounded by brackets. Outside and superscript to the brackets is a two negative sign.


(b) N2F4:

This Lewis structure shows two nitrogen atoms, each with one lone pair of electrons, single bonded to one another and each single bonded to two fluorine atoms. Each fluorine atom has three lone pairs of electrons.


(c) NH2:NH2:

This Lewis structure shows a nitrogen atom with two lone pairs of electrons single bonded to two hydrogen atoms. The structure is surrounded by brackets. Outside and superscript to the brackets is a negative sign.


(d) NF3:

This Lewis structure shows a nitrogen atom, with one lone pair of electrons, single bonded to three fluorine atoms. Each fluorine atom has three lone pairs of electrons.


(e) N3:N3:

Three Lewis structures are shown and connected by double-headed arrows in between. The left structure shows a nitrogen atom with a lone pair of electrons triple bonded to a second nitrogen which is single bonded to a third nitrogen. The third nitrogen has three lone pairs of electrons. The entire structure is surrounded by brackets, and outside and superscript to the brackets is a negative sign. The middle structure shows a nitrogen atom with three lone pair of electrons single bonded to a second nitrogen which is triple bonded to a third nitrogen. The third nitrogen which has one lone pair of electrons. The entire structure is surrounded by brackets, and outside and superscript to the brackets is a negative sign. The right structure shows a nitrogen atom with two lone pairs of electrons double bonded to a second nitrogen which is double bonded to a third nitrogen. The third nitrogen atom has two lone pairs of electrons. The entire structure is surrounded by brackets, and outside and superscript to the brackets is a negative sign.
61.

Ammonia acts as a Brønsted base because it readily accepts protons and as a Lewis base in that it has an electron pair to donate.
Brønsted base: NH3+H3O+NH4++H2ONH3+H3O+NH4++H2O
Lewis base: 2NH3+Ag+[H3NAgNH3]+2NH3+Ag+[H3NAgNH3]+

63.

(a) NO2:

Two Lewis structures are shown and connected by double-headed arrows in between. The left structure shows a nitrogen atom with a single electron double bonded to an oxygen atom which has two lone pairs of electrons. The nitrogen atom is also single bonded to an oxygen atom with three lone pairs of electrons. The right structure is a mirror image of the left structure.


Nitrogen is sp2 hybridized. The molecule has a bent geometry with an ONO bond angle of approximately 120°.
(b) NO2:NO2:

Two Lewis structures are shown and connected by double-headed arrows in between. Each structure is surrounded by brackets, and outside and superscript to the brackets is a negative sign. The left structure shows a nitrogen atom with a lone pair of electrons double bonded to an oxygen atom which has two lone pairs of electrons. The nitrogen atom is also single bonded to an oxygen atom with three lone pair of electrons. The right structure is a mirror image of the left structure.


Nitrogen is sp2 hybridized. The molecule has a bent geometry with an ONO bond angle slightly less than 120°.
(c) NO2+:NO2+:

This Lewis structure shows a nitrogen atom double bonded on both sides to an oxygen atom which has two lone pairs of electrons each. The structure is surrounded by brackets and outside and superscript to the brackets is a negative sign.


Nitrogen is sp hybridized. The molecule has a linear geometry with an ONO bond angle of 180°.

65.

Nitrogen cannot form a NF5 molecule because it does not have d orbitals to bond with the additional two fluorine atoms.

67.

(a)

This Lewis structure shows a phosphorus atom with a lone pair of electrons single bonded to three hydrogen atoms.


(b)

This Lewis structure shows a phosphorus atom single bonded to four hydrogen atoms. The structure is surrounded by brackets and has a superscript positive sign outside the brackets.


(c)

This Lewis structure shows two phosphorus atoms, each with a lone pair of electrons, single bonded to one another. Each phosphorus atom is also single bonded to two hydrogen atoms.


(d)

This Lewis structure shows a phosphorus atom single bonded to four oxygen atoms, each with three lone pairs of electrons. The structure is surrounded by brackets and has a superscript 3 negative sign outside the brackets.


(e)

This Lewis structure shows a phosphorus atom single bonded to five fluorine atoms, each with three lone pairs of electrons.
69.

(a) P4(s)+4Al(s)4AlP(s);P4(s)+4Al(s)4AlP(s); (b) P4(s)+12Na(s)4Na3P(s);P4(s)+12Na(s)4Na3P(s); (c) P4(s)+10F2(g)4PF5(l);P4(s)+10F2(g)4PF5(l); (d) P4(s)+6Cl2(g)4PCl3(l)P4(s)+6Cl2(g)4PCl3(l) or P4(s)+10Cl2(g)4PCl5(l);P4(s)+10Cl2(g)4PCl5(l); (e) P4(s)+3O2(g)P4O6(s)P4(s)+3O2(g)P4O6(s) or P4(s)+5O2(g)P4O10(s);P4(s)+5O2(g)P4O10(s); (f) P4O6(s)+2O2(g)P4O10(s)P4O6(s)+2O2(g)P4O10(s)

71.

291 mL

73.

28 tons

75.

(a)

This Lewis structure shows a phosphorus atom single bonded to four fluorine atoms, each with three lone pairs of electrons. The structure is surrounded by brackets and has a superscript positive sign outside the brackets. The label, “Tetrahedral,” is written under the structure.


(b)

This Lewis structure shows a phosphorus atom single bonded to five fluorine atoms, each with three lone pairs of electrons. The label, “Trigonal bipyramidal,” is written under the structure.


(c)

A Lewis structure shows a phosphorus atom single bonded to six fluorine atoms, each with three lone pairs of electrons. The structure is surrounded by brackets and has a superscript negative sign outside the brackets. The label, “Octahedral,” is written under the structure.


(d)

This Lewis structure shows a phosphorus atom single bonded to three fluorine atoms, each with three lone pairs of electrons. The phosphorus atom is also double bonded to an oxygen atom with two lone pairs of electrons. The label, “Tetrahedral,” is written under the structure.
77.

(a) P = 3+; (b) P = 5+; (c) P = 3+; (d) P = 5+; (e) P = 3−; (f) P = 5+

79.

FrO2

81.

(a) 2Zn(s)+O2(g)2ZnO(s);2Zn(s)+O2(g)2ZnO(s); (b) ZnCO3(s)ZnO(s)+CO2(g);ZnCO3(s)ZnO(s)+CO2(g); (c) ZnCO3(s)+2CH3COOH(aq)Zn(CH3COO)2(aq)+CO2(g)+H2O(l);ZnCO3(s)+2CH3COOH(aq)Zn(CH3COO)2(aq)+CO2(g)+H2O(l); (d) Zn(s)+2HBr(aq)ZnBr2(aq)+H2(g)Zn(s)+2HBr(aq)ZnBr2(aq)+H2(g)

83.

Al(OH)3(s)+3H+(aq)Al3++3H2O(l);Al(OH)3(s)+3H+(aq)Al3++3H2O(l); Al(OH)3(s)+OH[Al(OH)4](aq)Al(OH)3(s)+OH[Al(OH)4](aq)

85.

(a) Na2O(s)+H2O(l)2NaOH(aq);Na2O(s)+H2O(l)2NaOH(aq); (b) Cs2CO3(s)+2HF(aq)2CsF(aq)+CO2(g)+H2O(l);Cs2CO3(s)+2HF(aq)2CsF(aq)+CO2(g)+H2O(l); (c) Al2O3(s)+6HClO4(aq)2Al(ClO4)3(aq)+3H2O(l);Al2O3(s)+6HClO4(aq)2Al(ClO4)3(aq)+3H2O(l); (d) Na2CO3(aq)+Ba(NO3)2(aq)2NaNO3(aq)+BaCO3(s);Na2CO3(aq)+Ba(NO3)2(aq)2NaNO3(aq)+BaCO3(s); (e) TiCl4(l)+4Na(s)Ti(s)+4NaCl(s)TiCl4(l)+4Na(s)Ti(s)+4NaCl(s)

87.

HClO4 is the stronger acid because, in a series of oxyacids with similar formulas, the higher the electronegativity of the central atom, the stronger is the attraction of the central atom for the electrons of the oxygen(s). The stronger attraction of the oxygen electron results in a stronger attraction of oxygen for the electrons in the O-H bond, making the hydrogen more easily released. The weaker this bond, the stronger the acid.

89.

As H2SO4 and H2SeO4 are both oxyacids and their central atoms both have the same oxidation number, the acid strength depends on the relative electronegativity of the central atom. As sulfur is more electronegative than selenium, H2SO4 is the stronger acid.

91.

SO2, sp2 4+; SO3, sp2, 6+; H2SO4, sp3, 6+

93.

SF6: S = 6+; SO2F2: S = 6+; KHS: S = 2−

95.

Sulfur is able to form double bonds only at high temperatures (substantially endothermic conditions), which is not the case for oxygen.

97.

There are many possible answers including: Cu(s)+2H2SO4(l)CuSO4(aq)+SO2(g)+2H2O(l)Cu(s)+2H2SO4(l)CuSO4(aq)+SO2(g)+2H2O(l) and C(s)+2H2SO4(l)CO2(g)+2SO2(g)+2H2O(l)C(s)+2H2SO4(l)CO2(g)+2SO2(g)+2H2O(l)

99.

5.1 ×× 104 g

101.

SnCl4 is not a salt because it is covalently bonded. A salt must have ionic bonds.

103.

In oxyacids with similar formulas, the acid strength increases as the electronegativity of the central atom increases. HClO3 is stronger than HBrO3; Cl is more electronegative than Br.

105.

(a)

This Lewis structure shows an iodine atom with one lone pair of electrons single bonded to five fluorine atoms, each of which has three lone pairs of electrons. The image is labeled, “Square pyramidal.”


(b)

This Lewis structure shows an iodine atom with three lone pairs of electrons single bonded to two iodine atoms, each of which has three lone pairs of electrons. The image is surrounded by brackets. A superscript negative sign appears outside the brackets. The image is labeled, “Linear.”


(c)

This Lewis structure shows a phosphorus atom single bonded to five chlorine atoms, each of which has three lone pairs of electrons. The image is labeled, “Trigonal bipyramidal.”


(d)

This Lewis structure shows a selenium atom with one lone pair of electrons single bonded to four fluorine atoms, each of which has three lone pairs of electrons. The image is labeled “Seesaw.”


(e)

This Lewis structure shows a chlorine atom with two lone pairs of electrons single bonded to three fluorine atoms, each of which has three lone pairs of electrons. The image is labeled, “T-shaped.”
107.

(a) bromine trifluoride; (b) sodium bromate; (c) phosphorus pentabromide; (d) sodium perchlorate; (e) potassium hypochlorite

109.

(a) I: 7+; (b) I: 7+; (c) Cl: 4+; (d) I: 3+; Cl: 1−; (e) F: 0

111.

(a) sp3d hybridized; (b) sp3d2 hybridized; (c) sp3 hybridized; (d) sp3 hybridized; (e) sp3d2 hybridized;

113.

(a) nonpolar; (b) nonpolar; (c) polar; (d) nonpolar; (e) polar

115.

The empirical formula is XeF6, and the balanced reactions are: Xe(g)+3F2(g)ΔXeF6(s)XeF6(s)+3H2(g)6HF(g)+Xe(g)Xe(g)+3F2(g)ΔXeF6(s)XeF6(s)+3H2(g)6HF(g)+Xe(g)

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