Principles of General Chemistry

1

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Oracle Solaris 11 Advanced Administration Cookbook Packt Publishing Alexandre Borges Year: 2014 File: EPUB, 1.03 MB

Environmental science is grounded in chemistry. As one of many examples (discussed in Chapter 16),
a severe reduction in stratospheric ozone—an ozone “hole”—was confirmed over Antarctica in 1985.
This thin, 15-mile-high layer rich in ozone (O3, depicted as three connected red spheres) screens out
harmful solar ultraviolet (UV) light from reaching Earth’s surface, but Nobel-Prize winning research
showed that man-made chemicals were breaking O3 apart. In a series of reaction steps, Freon-12 (CCl2F2,
black sphere surrounded by two green and two yellow), escaping from air conditioners and spray cans,
rises intact through the air until it reaches the stratosphere. There, UV light splits off a chlorine atom
(Cl, green), which collides with an O3 molecule to form chlorine monoxide (ClO, red-green) and oxygen
(O2, two red). Regenerated in a later step, the Cl can then attack another O3. With a “lifetime” of about
2 years, each Cl atom can react with over 100,000 ozone molecules. To solve this problem, Freon-12 has
now been banned, and fewer Cl atoms have been detected in the ozone layer.

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Martin S. Silberberg
Third Edition

Principles of

GENERAL
CHEMISTRY

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PRINCIPLES OF GENERAL CHEMISTRY, THIRD EDITION
Published by McGraw-Hill, a business unit of The McGraw-Hill Companies, Inc., 1221 Avenue of the Americas,
New York, NY 10020. Copyright © 2013 by The McGraw-Hill Companies, Inc. All rights reserved. Printed in
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United States.
This book is printed on acid-free paper.
1 2 3 4 5 6 7 8 9 0 DOW/DOW 1 0 9 8 7 6 5 4 3 2
ISBN 978–0–07–340269–7
MHID 0–07–340269–9
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Library of Congress Cataloging-in-Publication Data
Silberberg, Martin S. (Martin Stuart), 1945Principles of general chemistry / Martin S. Silberberg. — 3rd ed.
p. cm.
Includes index.
ISBN 978–0–07–340269–7 — ISBN 0–07–340269–9 (hard copy : alk. paper) 1. Chemistry—Textbooks. I.
Title.
QD31.3.S55 2013
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2011015577

www.mhhe.com

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To Ruth and Daniel, with all my love
and
To the memory of my brother Bruce,
whose love, humor, and encouragement was
invaluable and will be profoundly missed.

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BRIEF CONTENTS
About the Author xvii
Preface xviii

1

Keys to the Study of Chemistry 2

2

The Components of Matter 32

3

Stoichiometry of Formulas and Equations 71

4

Three Major Classes of Chemical Reactions 115

5

Gases and the Kinetic-Molecular Theory 148

6

Thermochemistry: Energy Flow and Chemical Change 188

7

Quantum Theory and Atomic Structure 216

8

Electron Configuration and Chemical Periodicity 245

9

Models of Chemical Bonding 276

10 The Shapes of Molecules 302
11 Theories of Covalent Bonding 328
12 Intermolecular Forces: Liquids, Solids, and Phase Changes 350
13 The Properties of Solutions 391
14 Periodic Patterns in the Main-Group Elements 425
15 Organic Compounds and the Atomic Properties of Carbon 459
16 Kinetics: Rates and Mechanisms of Chemical Reactions 498
17 Equilibrium: The Extent of Chemical Reactions 542
18 Acid-Base Equilibria 579
19 Ionic Equilibria in Aqueous Systems 617
20 Thermodynamics: Entropy, Free Energy, and the Direction of Chemical Reactions 653
21 Electrochemistry: Chemical Change and Electrical Work 687
22 Transition Elements and Their Coordination Compounds 736
23 Nuclear Reactions and Their Applications 763
Appendix A Common Mathematical Operations in Chemistry A-1
Appendix B Standard Thermodynamic Values for Selected Substances A-5
Appendix C Equilibrium Constants for Selected Substances A-8
Appendix D Standard Electrode (Half-Cell) Potentials A-14
Appendix E Answers to Selected Problems A-15
Glossary G-1
Credits C-1
Index I-1

v

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DETAILED CONTENTS

CHAPTER 1

• Keys to the Study of Chemistry 2

1.1 Some Fundamental Definitions 3

1.2

The Properties of Matter 3
The States of Matter 5
The Central Theme in Chemistry 6
The Importance of Energy in the Study
of Matter 7
The Scientific Approach: Developing
a Model 9

1.3 Chemical Problem Solving 10

1.4

Units and Conversion Factors
in Calculations 10
A Systematic Approach to Solving
Chemistry Problems 12
Measurement in Scientific Study 14
General Features of SI Units 14
Some Important SI Units in Chemistry 14
Extensive and Intensive Properties 20

1.5 Uncertainty in Measurement:
Significant Figures 21
Determining Which Digits Are
Significant 21
Significant Figures: Calculations
and Rounding Off 22
Precision, Accuracy, and Instrument
Calibration 24
Chapter Review Guide 25
Problems 27

CHAPTER 2

• The Components of Matter   32

2.1 Elements, Compounds, and Mixtures:
2.2

2.3

2.4

An Atomic Overview 33
The Observations That Led to an
Atomic View of Matter 35
Mass Conservation 35
Definite Composition 36
Multiple Proportions 37
Dalton’s Atomic Theory 37
Postulates of the Atomic Theory 38
How the Theory Explains the Mass
Laws 38
The Observations That Led to the
Nuclear Atom Model 39
Discovery of the Electron and Its
Properties 39
Discovery of the Atomic Nucleus 41

2.5 The Atomic Theory Today 42

2.6
2.7

2.8

Structure of the Atom 42
Atomic Number, Mass Number,
and Atomic Symbol 43
Isotopes 44
Atomic Masses of the Elements;
Mass Spectrometry 44
Elements: A First Look at the Periodic
Table 47
Compounds: Introduction
to Bonding 49
The Formation of Ionic Compounds 49
The Formation of Covalent Compounds 51
Formulas, Names, and Masses
of Compounds 53
Binary Ionic Compounds 53

2.9

Compounds That Contain Polyatomic
Ions 56
Acid Names from Anion Names 57
Binary Covalent Compounds 58
The Simplest Organic Compounds:
Straight-Chain Alkanes 58
Molecular Masses from Chemical
Formulas 59
Representing Molecules with Formulas
and Models 60
Classification of Mixtures 61
An Overview of the Components
of Matter 62

Chapter Review Guide 63
Problems 65

vii

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viii

Detailed Contents

CHAPTER 3

• Stoichiometry of Formulas and Equations   71

3.1 The Mole 72

3.2

Defining the Mole 72
Determining Molar Mass 73
Converting Between Amount, Mass,
and Number of Chemical Entities 74
The Importance of Mass Percent 77
Determining the Formula
of an Unknown Compound 80
Empirical Formulas 80
Molecular Formulas 81
Isomers 84

CHAPTER 4
4.3

The Polar Nature of Water 116
Ionic Compounds in Water 116
Covalent Compounds in Water 120
Writing Equations for Aqueous Ionic
Reactions 120
Precipitation Reactions 122
The Key Event: Formation of a Solid from
Dissolved Ions 122
Predicting Whether a Precipitate Will
Form 123

CHAPTER 5

3.4

5.3

3.5 Fundamentals of Solution
Stoichiometry 99
Expressing Concentration in Terms
of Molarity 99
Amount-Mass-Number Conversions
Involving Solutions 100
Diluting a Solution 100
Stoichiometry of Reactions in Solution 103
Chapter Review Guide 105
Problems 108

4.4 Acid-Base Reactions 125

4.5

The Key Event: Formation of H2O from H1
and OH– 128
Proton Transfer in Acid-Base Reactions 129
Quantifying Acid-Base Reactions by
Titration 130
Oxidation-Reduction (Redox)
Reactions 132
The Key Event: Net Movement of Electrons
Between Reactants 132
Some Essential Redox Terminology 133

4.6

Using Oxidation Numbers to Monitor
Electron Charge 133
Elements in Redox Reactions 136
Combination Redox Reactions 136
Decomposition Redox Reactions 137
Displacement Redox Reactions and Activity
Series 137
Combustion Reactions 139

Chapter Review Guide 141
Problems 142

• Gases and the Kinetic-Molecular Theory   148

5.1 An Overview of the Physical States
5.2

Equations 85
Calculating Quantities of Reactant
and Product 89
Stoichiometrically Equivalent Molar Ratios
from the Balanced Equation 89
Reactions That Involve a Limiting
Reactant 93
Theoretical, Actual, and Percent Reaction
Yields 97

• Three Major Classes of Chemical Reactions   115

4.1 The Role of Water as a Solvent 116

4.2

3.3 Writing and Balancing Chemical

of Matter 149
Gas Pressure and Its
Measurement 150
Measuring Atmospheric Pressure 150
Units of Pressure 151
The Gas Laws and Their Experimental
Foundations 153
The Relationship Between Volume and
Pressure: Boyle’s Law 153
The Relationship Between Volume and
Temperature: Charles’s Law 154

5.4

The Relationship Between Volume and
Amount: Avogadro’s Law 156
Gas Behavior at Standard Conditions 156
The Ideal Gas Law 157
Solving Gas Law Problems 158
Rearrangements of the Ideal
Gas Law 162
The Density of a Gas 162
The Molar Mass of a Gas 164
The Partial Pressure of a Gas in a Mixture
of Gases 165
The Ideal Gas Law and Reaction
Stoichiometry 167

5.5 The Kinetic-Molecular Theory:

5.6

A Model for Gas Behavior 170
How the Kinetic-Molecular Theory
Explains the Gas Laws 170
Effusion and Diffusion 175
Real Gases: Deviations from Ideal
Behavior 177
Effects of Extreme Conditions on Gas
Behavior 177
The van der Waals Equation: Adjusting
the Ideal Gas Law 179

Chapter Review Guide 179
Problems 182

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ix

CHAPTER 6

• Thermochemistry: Energy Flow and Chemical Change   188

6.1 Forms of Energy and Their
Interconversion 189
Defining the System and Its
Surroundings 189
Energy Transfer to and from a System 190
Heat and Work: Two Forms of Energy
Transfer 190
The Law of Energy Conservation 192
Units of Energy 193
State Functions and the Path Independence
of the Energy Change 194

CHAPTER 7

The Wave Nature of Light 217
The Particle Nature of Light 220
Atomic Spectra 223
Line Spectra and the Rydberg Equation 223
The Bohr Model of the Hydrogen Atom 224
The Energy Levels of the Hydrogen
Atom 226
Spectral Analysis in the Laboratory 228

CHAPTER 8

Atoms 246
The Electron-Spin Quantum Number 246
The Exclusion Principle and Orbital
Occupancy 247
Electrostatic Effects and Energy-Level
Splitting 247
The Quantum-Mechanical Model
and the Periodic Table 249
Building Up Period 1 249
Building Up Period 2 250

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6.4

6.5 Hess’s Law: Finding DH of Any
Reaction 203

6.6 Standard Enthalpies of Reaction
(DHrxn) 205
Formation Equations and Their Standard
Enthalpy Changes 205
Determining DH rxn from DH f Values for
Reactants and Products 206
Fossil Fuels and Climate Change 207
Chapter Review Guide 209
Problems 211

7.3 The Wave-Particle Duality of Matter

7.4

and Energy 229
The Wave Nature of Electrons and the
Particle Nature of Photons 229
Heisenberg’s Uncertainty Principle 231
The Quantum-Mechanical Model
of the Atom 232
The Atomic Orbital and the Probable
Location of the Electron 232

Quantum Numbers of an Atomic
Orbital 234
Quantum Numbers and Energy Levels 235
Shapes of Atomic Orbitals 237
The Special Case of Energy Levels
in the H Atom 239
Chapter Review Guide 240
Problems 241

• Electron Configuration and Chemical Periodicity   245

8.1 Characteristics of Many-Electron

8.2

6.3

Constant Pressure 195
The Meaning of Enthalpy 195
Exothermic and Endothermic Processes 196
Calorimetry: Measuring the Heat of a
Chemical or Physical Change 197
Specific Heat Capacity 197
The Two Common Types
of Calorimetry 198
Stoichiometry of Thermochemical
Equations 201

• Quantum Theory and Atomic Structure   216

7.1 The Nature of Light 217
7.2

6.2 Enthalpy: Chemical Change at

Building Up Period 3 251
Similar Electron Configurations Within
Groups 252
Building Up Period 4: The First Transition
Series 253
General Principles of Electron
Configurations 254
Intervening Series: Transition and Inner
Transition Elements 256

8.3 Trends in Three Atomic

8.4

Properties 258
Trends in Atomic Size 258
Trends in Ionization Energy 260
Trends in Electron Affinity 263
Atomic Properties and Chemical
Reactivity 265
Trends in Metallic Behavior 265
Properties of Monatomic Ions 266

Chapter Review Guide 271
Problems 272

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x

Detailed Contents

CHAPTER 9

• Models of Chemical Bonding   276

9.1 Atomic Properties and Chemical

9.2

9.3

Bonds 277
Types of Bonding: Three Ways Metals
and Nonmetals Combine 277
Lewis Symbols and the Octet Rule 278
The Ionic Bonding Model 280
Why Ionic Compounds Form: The
Importance of Lattice Energy 280
Periodic Trends in Lattice Energy 281
How the Model Explains the Properties of
Ionic Compounds 283
The Covalent Bonding Model 284
The Formation of a Covalent Bond 284
Bonding Pairs and Lone Pairs 285

CHAPTER 10

Lewis Structures 303
Applying the Octet Rule to Write Lewis
Structures 303
Resonance: Delocalized Electron-Pair
Bonding 306
Formal Charge: Selecting the More
Important Resonance Structure 308
Lewis Structures for Exceptions to the Octet
Rule 309
Valence-Shell Electron-Pair Repulsion
(VSEPR) Theory and Molecular
Shape 312
Electron-Group Arrangements and
Molecular Shapes 312

CHAPTER 11

Electronegativity and Bond
Polarity 293
Electronegativity 293
Bond Polarity and Partial Ionic
Character 294
The Gradation in Bonding Across
a Period 296
Chapter Review Guide 297
Problems 298

The Molecular Shape with Two Electron
Groups (Linear Arrangement) 313
Molecular Shapes with Three
Electron Groups (Trigonal Planar
Arrangement) 314
Molecular Shapes with Four Electron
Groups (Tetrahedral Arrangement) 314
Molecular Shapes with Five Electron
Groups (Trigonal Bipyramidal
Arrangement) 316
Molecular Shapes with Six Electron Groups
(Octahedral Arrangement) 317

10.3

Using VSEPR Theory to Determine
Molecular Shape 318
Molecular Shapes with More Than One
Central Atom 319
Molecular Shape and Molecular
Polarity 320
Bond Polarity, Bond Angle, and Dipole
Moment 321

Chapter Review Guide 322
Problems 324

• Theories of Covalent Bonding   328

11.1 Valence Bond (VB) Theory and Orbital
Hybridization 329
The Central Themes of VB
Theory 329
Types of Hybrid Orbitals 330

9.5 Between the Extremes:

• The Shapes of Molecules   302

10.1 Depicting Molecules and Ions with

10.2

9.4

Properties of a Covalent Bond: Order,
Energy, and Length 285
How the Model Explains the Properties
of Covalent Substances 288
Using IR Spectroscopy to Study Covalent
Compounds 289
Bond Energy and Chemical
Change 290
Changes in Bond Energy: Where Does
DH rxn Come From? 290
Using Bond Energies to Calculate
DH rxn 290

11.2 Modes of Orbital Overlap and the

11.3 Molecular Orbital (MO) Theory and

Types of Covalent Bonds 335
Orbital Overlap in Single and Multiple
Bonds 335
Orbital Overlap and Molecular
Rotation 337

Electron Delocalization 338
The Central Themes of MO Theory 338
Homonuclear Diatomic Molecules
of Period 2 Elements 341
Chapter Review Guide 345
Problems 346

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xi

CHAPTER 12

• Intermolecular Forces: Liquids, Solids, and Phase Changes   350

12.1 An Overview of Physical States and
12.2

12.3

Phase Changes 351
Quantitative Aspects of Phase
Changes 354
Heat Involved in Phase Changes 354
The Equilibrium Nature of Phase
Changes 357
Phase Diagrams: Effect of Pressure and
Temperature on Physical State 360
Types of Intermolecular Forces 362
How Close Can Molecules Approach Each
Other? 362
Ion-Dipole Forces 362
Dipole-Dipole Forces 363

CHAPTER 13

12.4

12.5

Forces and Solubility 392
Intermolecular Forces in Solution 393
Liquid Solutions and the Role of Molecular
Polarity 394
Gas Solutions and Solid Solutions 396
Why Substances Dissolve:
Understanding the Solution
Process 397
Heats of Solution: Solution Cycles 397

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12.6 The Solid State: Structure, Properties,
and Bonding 373
Structural Features of Solids 373
Types and Properties of Crystalline
Solids 379
Bonding in Solids I: The Electron-Sea
Model of Metallic Bonding 382
Bonding in Solids II: Band Theory 382
Chapter Review Guide 385
Problems 386

• The Properties of Solutions   391

13.1 Types of Solutions: Intermolecular

13.2

The Hydrogen Bond 364
Polarizability and Induced Dipole
Forces 366
Dispersion (London) Forces 366
Properties of the Liquid State 369
Surface Tension 369
Capillarity 370
Viscosity 370
The Uniqueness of Water 371
Solvent Properties of Water 371
Thermal Properties of Water 371
Surface Properties of Water 372
The Unusual Density of Solid Water 372

13.3

13.4

Heats of Hydration: Ionic Solids
in Water 397
The Solution Process and the Change
in Entropy 399
Solubility as an Equilibrium
Process 401
Effect of Temperature on Solubility 401
Effect of Pressure on Solubility 402
Concentration Terms 404
Molarity and Molality 404
Parts of Solute by Parts of Solution 405
Interconverting Concentration Terms 406

13.5 Colligative Properties
of Solutions 408
Nonvolatile Nonelectrolyte Solutions 408
Using Colligative Properties to Find Solute
Molar Mass 413
Volatile Nonelectrolyte Solutions 414
Strong Electrolyte Solutions 415
Chapter Review Guide 417
Problems 419

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xii

Detailed Contents

CHAPTER 14

• Periodic Patterns in the Main-Group Elements   425

14.1 Hydrogen, the Simplest Atom 426

14.2

14.3

14.4

Where Hydrogen Fits in the Periodic
Table 426
Highlights of Hydrogen Chemistry 426
Group 1A(1): The Alkali Metals 427
Why the Alkali Metals Have Unusual
Physical Properties 427
Why the Alkali Metals Are So Reactive 427
The Anomalous Behavior of Period 2
Members 429
Group 2A(2): The Alkaline Earth
Metals 430
How the Alkaline Earth and Alkali Metals
Compare Physically 430
How the Alkaline Earth and Alkali Metals
Compare Chemically 430
Diagonal Relationships 432
Group 3A(13): The Boron Family 432
How the Transition Elements Influence
Properties 432

CHAPTER 15

14.5

14.6

14.7

and the Characteristics of Organic
Molecules 460
The Structural Complexity of Organic
Molecules 460
The Chemical Diversity of Organic
Molecules 461
The Structures and Classes of
Hydrocarbons 462
Carbon Skeletons and Hydrogen Skins 462
Alkanes: Hydrocarbons with Only Single
Bonds 464
Constitutional Isomerism and the Physical
Properties of Alkanes 467
Chiral Molecules and Optical
Isomerism 468

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14.8

14.9

How the Oxygen and Nitrogen Families
Compare Chemically 446
Highlights of Oxygen Chemistry 447
Highlights of Sulfur Chemistry 447
Group 7A(17): The Halogens 448
How the Halogens and the Alkali Metals
Contrast Physically 448
Why the Halogens Are So Reactive 448
Highlights of Halogen Chemistry 450
Group 8A(18): The Noble Gases 452
Physical Properties 452
Why Noble Gases Can Form
Compounds 452

Chapter Review Guide 452
Problems 454

• Organic Compounds and the Atomic Properties of Carbon  459

15.1 The Special Nature of Carbon

15.2

Features That First Appear in This Group’s
Chemical Properties 432
Group 4A(14): The Carbon Family 434
How the Type of Bonding in an Element
Affects Physical Properties 434
How Bonding Changes in the Carbon
Family’s Compounds 436
Highlights of Carbon Chemistry 436
Highlights of Silicon Chemistry 438
Group 5A(15): The Nitrogen
Family 439
The Wide Range of Physical Behavior 439
Patterns in Chemical Behavior 439
Highlights of Nitrogen Chemistry 441
Highlights of Phosphorus Chemistry 443
Group 6A(16): The Oxygen
Family 444
How the Oxygen and Nitrogen Families
Compare Physically 446

15.3
15.4

Alkenes: Hydrocarbons with
Double Bonds 469
Alkynes: Hydrocarbons with Triple
Bonds 470
Aromatic Hydrocarbons: Cyclic Molecules
with Delocalized π Electrons 471
Some Important Classes of Organic
Reactions 472
Properties and Reactivities of
Common Functional Groups 474
Functional Groups with Only Single
Bonds 474
Functional Groups with Double Bonds 478
Functional Groups with Both Single and
Double Bonds 480
Functional Groups with Triple Bonds 482

15.5 The Monomer-Polymer Theme I:

15.6

Synthetic Macromolecules 483
Addition Polymers 484
Condensation Polymers 484
The Monomer-Polymer Theme II:
Biological Macromolecules 486
Sugars and Polysaccharides 486
Amino Acids and Proteins 487
Nucleotides and Nucleic Acids 490

Chapter Review Guide 492
Problems 494

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xiii

CHAPTER 16

• Kinetics: Rates and Mechanisms of Chemical Reactions   498

16.1 Focusing on Reaction Rate 499
16.2 Expressing the Reaction Rate 501

16.3

16.4

Average, Instantaneous, and Initial Reaction
Rates 501
Expressing Rate in Terms of Reactant and
Product Concentrations 503
The Rate Law and Its
Components 505
Some Laboratory Methods for Determining
the Initial Rate 505
Determining Reaction Orders 505
Determining the Rate Constant 512
Integrated Rate Laws: Concentration
Changes Over Time 512
Integrated Rate Laws for First-, Second-,
and Zero-Order Reactions 513

CHAPTER 17
17.3
17.4

Equilibrium Constant 543
The Reaction Quotient and the
Equilibrium Constant 545
Changing Value of the Reaction
Quotient 545
Writing the Reaction Quotient 546
Expressing Equilibria with Pressure
Terms: Relation Between Kc
and Kp 550
Comparing Q and K to Predict
Reaction Direction 552

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16.6

16.7 Catalysis: Speeding Up
a Reaction 530
The Basis of Catalytic Action 531
Homogeneous Catalysis 531
Heterogeneous Catalysis 532
Catalysis in Nature 533
Chapter Review Guide 534
Problems 536

• Equilibrium: The Extent of Chemical Reactions   542

17.1 The Equilibrium State and the
17.2

16.5

Determining Reaction Orders from an
Integrated Rate Law 514
Reaction Half-Life 515
Theories of Chemical Kinetics 519
Collision Theory: Basis of the
Rate Law 519
Transition State Theory: What the
Activation Energy Is Used For 522
Reaction Mechanisms: The Steps from
Reactant to Product 525
Elementary Reactions and Molecularity 526
The Rate-Determining Step of a Reaction
Mechanism 527
Correlating the Mechanism with the Rate
Law 528

17.5 How to Solve Equilibrium

17.6

Problems 554
Using Quantities to Find the Equilibrium
Constant 554
Using the Equilibrium Constant to Find
Quantities 556
Problems Involving Mixtures of Reactants
and Products 561
Reaction Conditions and Equilibrium:
Le Châtelier’s Principle 562
The Effect of a Change in
Concentration 563

The Effect of a Change in Pressure
(Volume) 565
The Effect of a Change in Temperature 567
The Lack of Effect of a Catalyst 568
The Industrial Production of Ammonia 570
Chapter Review Guide 572
Problems 573

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xiv

Detailed Contents

CHAPTER 18

• Acid-Base Equilibria   579

18.1 Acids and Bases in Water 580

18.2

18.3

Release of H1 or OH2 and the Arrhenius
Acid-Base Definition 580
Variation in Acid Strength: The AcidDissociation Constant (Ka) 581
Classifying the Relative Strengths of Acids
and Bases 583
Autoionization of Water and the pH
Scale 584
The Equilibrium Nature of Autoionization:
The Ion-Product Constant for Water
(Kw) 584
Expressing the Hydronium Ion
Concentration: The pH Scale 585
Proton Transfer and the BrønstedLowry Acid-Base Definition 588
Conjugate Acid-Base Pairs 589
Relative Acid-Base Strength and the Net
Direction of Reaction 590

CHAPTER 19

18.4 Solving Problems Involving Weak-Acid

18.5

18.6

What a Buffer Is and How It Works: The
Common-Ion Effect 618
The Henderson-Hasselbalch Equation 622
Buffer Capacity and Buffer Range 623
Preparing a Buffer 625
Acid-Base Titration Curves 626
Monitoring pH with Acid-Base
Indicators 627
Strong Acid–Strong Base Titration
Curves 627

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18.7 Acid-Base Properties of Salt

18.8

Solutions 603
Salts That Yield Neutral Solutions 604
Salts That Yield Acidic Solutions 604
Salts That Yield Basic Solutions 604
Salts of Weakly Acidic Cations and Weakly
Basic Anions 605
Salts of Amphiprotic Anions 605
Electron-Pair Donation and the Lewis
Acid-Base Definition 607
Molecules as Lewis Acids 607
Metal Ions as Lewis Acids 608

Chapter Review Guide 609
Problems 611

• Ionic Equilibria in Aqueous Systems   617

19.1 Equilibria of Acid-Base Buffers 618

19.2

Equilibria 593
Finding Ka Given Concentrations 593
Finding Concentrations Given Ka 594
The Effect of Concentration on the Extent
of Acid Dissociation 595
The Behavior of Polyprotic Acids 597
Weak Bases and Their Relation to
Weak Acids 597
Molecules as Weak Bases: Ammonia and
the Amines 598
Anions of Weak Acids as Weak Bases 599
The Relation Between Ka and Kb of a
Conjugate Acid-Base Pair 600
Molecular Properties and Acid
Strength 601
Acid Strength of Nonmetal Hydrides 601
Acid Strength of Oxoacids 602
Acidity of Hydrated Metal Ions 602

19.3

Weak Acid–Strong Base Titration
Curves 629
Weak Base–Strong Acid Titration
Curves 632
Equilibria of Slightly Soluble Ionic
Compounds 633
The Ion-Product Expression (Qsp) and the
Solubility-Product Constant (Ksp) 634
Calculations Involving the SolubilityProduct Constant 635
Effect of a Common Ion on Solubility 637

19.4

Effect of pH on Solubility 639
Predicting the Formation of a Precipitate:
Qsp vs. Ksp 639
Ionic Equilibria and the Acid-Rain
Problem 641
Equilibria Involving Complex Ions 643
Formation of Complex Ions 643
Complex Ions and Solubility
of Precipitates 645

Chapter Review Guide 646
Problems 648

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xv

CHAPTER 20

• Thermodynamics: Entropy, Free Energy,
and the Direction of Chemical Reactions   653

20.1 The Second Law of Thermodynamics:
Predicting Spontaneous Change 654
The First Law of Thermodynamics Does Not
Predict Spontaneous Change 654
The Sign of DH Does Not Predict
Spontaneous Change 655
Freedom of Particle Motion and Dispersal of
Particle Energy 655
Entropy and the Number of Microstates 656
Entropy and the Second Law of
Thermodynamics 659
Standard Molar Entropies and the Third
Law 659
Predicting Relative S° of a System 660

CHAPTER 21

20.2 Calculating the Change in Entropy

20.3

21.3

Cells 688
A Quick Review of Oxidation-Reduction
Concepts 688
Half-Reaction Method for Balancing Redox
Reactions 689
An Overview of Electrochemical Cells 692
Voltaic Cells: Using Spontaneous
Reactions to Generate Electrical
Energy 693
Construction and Operation of a Voltaic
Cell 694
Notation for a Voltaic Cell 696
Cell Potential: Output of a Voltaic
Cell 697
Standard Cell Potentials 697
Relative Strengths of Oxidizing and
Reducing Agents 700

siL02699_fm_i_xxvii.indd 15

20.4

Calculating Standard Free Energy
Changes 669
The Free Energy Change and the Work a
System Can Do 671
The Effect of Temperature on Reaction
Spontaneity 671
Coupling of Reactions to Drive a
Nonspontaneous Change 674
Free Energy, Equilibrium, and
Reaction Direction 676

Chapter Review Guide 681
Problems 682

• Electrochemistry: Chemical Change and Electrical Work   687

21.1 Redox Reactions and Electrochemical

21.2

of a Reaction 664
Entropy Changes in the System: Standard
Entropy of Reaction (DSrxn) 664
Entropy Changes in the Surroundings: The
Other Part of the Total 665
The Entropy Change and the Equilibrium
State 667
Spontaneous Exothermic and Endothermic
Changes 667
Entropy, Free Energy, and Work 668
Free Energy Change and Reaction
Spontaneity 668

21.4

21.5

21.6

Writing Spontaneous Redox Reactions 701
Explaining the Activity Series
of the Metals 704
Free Energy and Electrical Work 705
Standard Cell Potential and the Equilibrium
Constant 705
The Effect of Concentration on Cell
Potential 707
Changes in Potential During Cell
Operation 709
Concentration Cells 710
Electrochemical Processes
in Batteries 713
Primary (Nonrechargeable) Batteries 713
Secondary (Rechargeable) Batteries 714
Fuel Cells 716
Corrosion: An Environmental Voltaic
Cell 717
The Corrosion of Iron 717

21.7

Protecting Against the Corrosion
of Iron 718
Electrolytic Cells: Using Electrical
Energy to Drive Nonspontaneous
Reactions 719
Construction and Operation of an
Electrolytic Cell 719
Predicting the Products of Electrolysis 720
Purifying Copper and Isolating
Aluminum 724
Stoichiometry of Electrolysis: The Relation
Between Amounts of Charge and
Products 726

Chapter Review Guide 728
Problems 730

11/29/11 10:25 AM

xvi

Detailed Contents

CHAPTER 22

• Transition Elements and Their Coordination Compounds   736

22.1 Properties of the Transition

22.2 Coordination Compounds 743

Elements 737
Electron Configurations of the Transition
Metals and Their Ions 738
Atomic and Physical Properties of the
Transition Elements 739
Chemical Properties of the Transition
Elements 741

CHAPTER 23

23.3

Stability 764
The Components of the Nucleus: Terms
and Notation 764
Modes of Radioactive Decay; Balancing
Nuclear Equations 765
Nuclear Stability and the Mode
of Decay 768
The Kinetics of Radioactive Decay 772
The Rate of Radioactive Decay 772
Radioisotopic Dating 774
Nuclear Transmutation: Induced
Changes in Nuclei 776

Appendix A Common Mathematical
Operations in Chemistry A-1
Appendix B Standard Thermodynamic
Values for Selected Substances A-5
Appendix C Equilibrium Constants
for Selected Substances A-8

siL02699_fm_i_xxvii.indd 16

Crystal Field Theory 752
Transition Metal Complexes in Biological
Systems 757
Chapter Review Guide 758
Problems 759

• Nuclear Reactions and Their Applications   763

23.1 Radioactive Decay and Nuclear

23.2

22.3

Complex Ions: Coordination Numbers,
Geometries, and Ligands 744
Formulas and Names of Coordination
Compounds 745
Isomerism in Coordination Compounds 747
Theoretical Basis for the Bonding
and Properties of Complexes 750
Applying Valence Bond Theory to Complex
Ions 750

23.4 Effects of Nuclear Radiation

23.5

23.6

on Matter 778
Effects of Ionizing Radiation
on Living Tissue 778
Sources of Ionizing Radiation 779
Applications of Radioisotopes 780
Radioactive Tracers 780
Additional Applications of Ionizing
Radiation 782
The Interconversion of Mass
and Energy 783
The Mass Difference Between a Nucleus
and Its Nucleons 783

Appendix D Standard Electrode
(Half-Cell) Potentials A-14
Appendix E Answers to Selected
Problems A-15

23.7

Nuclear Binding Energy and the Binding
Energy per Nucleon 784
Applications of Fission and Fusion 786
The Process of Nuclear Fission 786
The Promise of Nuclear Fusion 789

Chapter Review Guide 790
Problems 792

Glossary G-1
Credits C-1
Index I-1

11/29/11 10:25 AM

About the Author
Martin S. Silberberg received a B.S. in Chemistry from the City University of New

York and a Ph.D. in Chemistry from the University of Oklahoma. He then accepted a
research position in analytical biochemistry at the Albert Einstein College of Medicine
in New York City, where he developed advanced methods to study fundamental brain
mechanisms as well as neurotransmitter metabolism in Parkinson’s disease. Following
his years in research, Dr. Silberberg joined the faculty of Bard College at Simon’s
Rock, a liberal arts college known for its excellence in teaching small classes of highly
motivated students. As Head of the Natural Sciences Major and Director of Premedical
Studies, he taught courses in general chemistry, organic chemistry, biochemistry, and
liberal arts chemistry. The close student contact afforded him insights into how students learn chemistry, where they have difficulties, and what strategies can help them
succeed. Prof. Silberberg applied these insights in a broader context by establishing a
text writing, editing, and consulting company. Before writing his own text, he worked
as a consulting and developmental editor on chemistry, biochemistry, and physics texts
for several major college publishers. He resides with his wife and son in the Pioneer
Valley near Amherst, Massachusetts, where he enjoys the rich cultural and academic
life of the area and relaxes by cooking, singing, and hiking.

xvii

siL02699_fm_i_xxvii.indd 17

11/29/11 10:25 AM

Preface
As the new century unfolds, chemistry will play its usual,
crucial role in dealing with complex environmental, medical, and industrial issues. And, as the complexities increase
and more information is needed to understand them, many
chemistry instructors want a more focused text to serve as
the core of a powerful electronic teaching and learning
package. This new, Third Edition of Principles of General
Chemistry is the ideal choice, designed to cover key principles and skills with great readability, the most accurate
molecular art available, a problem-solving approach that is
universally praised, and a supporting suite of electronic
products that sets a new standard in academic science.

How Principles and Chemistry
Are the Same
Principles of General Chemistry was created from its parent
text, Chemistry: The Molecular Nature of Matter and
Change, when four expert chemistry teachers—three consulting professors and the author—joined to distill the
concepts and skills at the heart of general chemistry. Principles covers all the material a science major needs to continue in premedical studies, engineering, or related fields.
It maintains the same high standards of accuracy, clarity,
and rigor as its parent and adopts the same three distinguishing hallmarks:
1. Visualizing chemical models. In many places in the text,
concepts are explained first at the macroscopic level and
then from a molecular point of view. Placed near many of
these discussions, the text’s celebrated graphics depict the
phenomenon or change at the observable level in the lab,
at the atomic level with superbly accurate molecular art,
and at the symbolic level with the balanced equation.
2. Thinking logically to solve problems. The problem-solving
approach, based on a four-step method widely approved
by chemical educators, is introduced in Chapter 1 and
employed consistently throughout the text. It encourages
students to first plan a logical approach, and only then
proceed to the arithmetic solution. A check step, universally recommended by instructors, fosters the habit of considering the reasonableness and magnitude of the answer.
For practice and reinforcement, each worked problem has
a matched follow-up problem, for which an abbreviated,
multistep solution—not merely a numerical answer—
appears at the end of the chapter.
3. Applying ideas to the real world. For today’s students, who
may enter one of numerous chemistry-related fields, especially important applications—such as climate change,

enzyme catalysis, materials science, and others—are
woven into the text discussion, and real-world scenarios
are used in many worked in-chapter sample problems as
well as end-of-chapter problems.

Principles and Chemistry also share a common topic
sequence, which provides a thorough introduction to chemistry for science majors:
• Chapters 1 through 6 cover unit conversions and uncertainty, introduce atomic structure and bonding, discuss
stoichiometry and reaction classes, show how gas behavior
is modeled, and highlight the relation between heat and
chemical change.
• Chapters 7 through 15 take an “atoms-first” approach, as
they move from atomic structure and electron configuration to how atoms bond and what the resulting molecules
look like and why. Intermolecular forces are covered by
discussing the behavior of liquids and solids as compared
with that of gases, and then leads the different behavior of
solutions. These principles are then applied to the chemistry of the elements and to the compounds of carbon.
• Chapters 16 through 21 cover dynamic aspects of reaction
chemistry, including kinetics, equilibrium, entropy and
free energy, and electrochemistry.
• Chapters 22 and 23 cover transition elements and nuclear
reactions.

How Principles and Chemistry
Are Different
Principles presents the same authoritative coverage as
Chemistry but in 240 fewer pages. It does so by removing
most of the boxed application material, thus letting instructors choose applications tailored for their course. Moreover,
several topics that are important areas of research but not
central to general chemistry were left out, including colloids, polymers, liquid crystals, and so forth. And mainstream material from the chapter on isolating the elements
was blended into the chapter on electrochemistry.
Despite its much shorter length, Principles of General
Chemistry includes all the pedagogy so admired in Chemistry. It has all the worked sample problems and about twothirds as many end-of-chapter problems, still more than
enough problems for every topic, with a high level of relevance and many real-world applications. The learning aids
that students find so useful have also been retained—
Concepts and Skills to Review, Section Summaries, Key
Terms, Key Equations, and Brief Solutions to Follow-up
Problems.

xviii

siL02699_fm_i_xxvii.indd 18

11/29/11 10:25 AM

Preface

In addition, three aids not found in the parent Chemistry
help students focus their efforts:
• Key Principles. At the beginning of each chapter, short
bulleted paragraphs state the main concepts concisely,
using many of the same phrases and terms (in italics)
that appear in the pages to follow. A student can preview
these principles before reading the chapter and then
review them afterward.
• “Think of It This Way . . .” with Analogies, Mnemonics,
and Insights. This recurring feature provides analogies
for difficult concepts (e.g., the “radial probability distribution” of apples around a tree) and amazing quantities
(e.g., a stadium and a marble for the relative sizes of
atom and nucleus), memory shortcuts (e.g., which reaction occurs at which electrode), and useful insights
(e.g., similarities between a saturated solution and a
liquid-vapor system).
• Problem-Based Learning Objectives. The list of learning
objectives at the end of each chapter includes the end-ofchapter problems that relate to each objective. Thus, a
student, or instructor, can select problems that review a
given topic.

What’s New in the Third Edition
To address dynamic changes in how courses are structured
and how students learn—variable math and reading preparation, less time for traditional studying, electronic media as
part of lectures and homework, new challenges and options
in career choices—the author and publisher consulted
extensively with students and faculty. Based on their input,
we developed the following ways to improve the text as a
whole as well as the content of individual chapters.

Global Changes to the Entire Text
Writing style and content presentation. Every line of
every discussion has been revised to optimize clarity, readability, and a more direct presentation. The use of additional
subheads, numbered (and titled) paragraphs, and bulleted
(and titled) lists has eliminated long unbroken paragraphs.
Main ideas are delineated and highlighted, making for more
efficient study and lectures. As a result, the text is over 20
pages shorter than the Second Edition.
More worked problems. The much admired—and imitated—four-part (plan, solution, check, practice) Sample
Problems occur in both data-based and molecular-scene
format. To deepen understanding, Follow-up Problems have
worked-out solutions at the back of each chapter, with a
road map when appropriate, effectively doubling the number of worked problems. This edition has 15 more sample
problems, many in the earlier chapters, where students need
the most practice in order to develop confidence.

siL02699_fm_i_xxvii.indd 19

xix

Art and figure legends. Figures have been made more
realistic and modern. Figure legends have been greatly
shortened, and the explanations from them have either been
added to the text or included within the figures.
Page design and layout. A more open look invites the
reader while maintaining the same attention to keeping text
and related figures and tables near each other for easier
studying.
Section summaries. This universally approved feature
is even easier to use in a new bulleted format.
Chapter review. The unique Chapter Review Guide aids
study with problem-based learning objectives, key terms,
key equations, and the multistep Brief Solutions to Followup Problems (rather than just numerical answers).
End-of-chapter problem sets. With an enhanced design
to improve readability and traditional and molecular-scene
problems updated and revised, these problem sets are far
more extensive than in other brief texts.

Content Changes to Individual Chapters
• Chapter 2 presents a new figure and table on molecular
modeling, and it addresses the new IUPAC recommendations for atomic masses.
• Discussion of empirical formulas has been moved from
Chapter 2 to Chapter 3 so that it appears just before
molecular formulas.
• Chapter 3 has some sample problems from the Second
Edition that have been divided to focus on distinct concepts, and it contains seven new sample problems.
• Chapters 3 and 4 include more extensive and consistent
use of stoichiometry reaction tables in limiting-reactant
problems.
• Chapter 4 presents a new molecular-scene sample problem on depicting an ionic compound in aqueous solution.
• Chapter 5 includes a new discussion on how gas laws
apply to breathing.
• Chapter 5 groups stoichiometry of gaseous reactions with
other rearrangements of the ideal gas law.
• Chapter 17 makes consistent use of quantitative benchmarks for determining when it is valid to assume that
the amount reacting can be neglected.

Acknowledgments
For the third edition of Principles of General Chemistry, I
am once again very fortunate that Patricia Amateis of Virginia Tech prepared the Instructors’ Solutions Manual and
Student Solutions Manual and Libby Weberg the Student
Study Guide.
The following individuals helped write and review
goal-oriented content for LearnSmart for general chemistry:
Erin Whitteck; Margaret Ruth Leslie, Kent State University;
and Adam I. Keller, Columbus State Community College.

11/29/11 10:25 AM

xx

Preface

And, I greatly appreciate the efforts
of all the professors who reviewed portions of the new edition or who participated in our developmental survey to
assess the content needs for the text:
DeeDee A. Allen, Wake Technical
Community College
John D. Anderson, Midland College
Jeanne C. Arquette, Phoenix College
Yiyan Bai, Houston Community College
Stanley A. Bajue, Medgar Evers College, CUNY
Jason P. Barbour, Anne Arundel Community
College
Peter T. Bell, Tarleton State University
Vladimir Benin, University of Dayton
Paul J. Birckbichler, Slippery Rock University
Simon Bott, University of Houston
Kevin A. Boudreaux, Angelo State University
R. D. Braun, University of Louisiana,
Lafayette
Stacey Buchanan, Henry Ford Community
College
Michael E. Clay, College of San Mateo
Michael Columbia, Indiana University
Purdue University Fort Wayne
Charles R. Cornett, University of Wisconsin,
Platteville
Kevin Crawford, The Citadel
Mapi M. Cuevas, Santa Fe Community
College
Kate Deline, College of San Mateo
Amy M. Deveau, University of New England,
Biddeford
Jozsef Devenyi, The University of Tennessee,
Martin
Paul A. DiMilla, Northeastern University
John P. DiVincenzo, Middle Tennessee State
University
Ajit Dixit, Wake Technical Community
College
Son Q. Do, University of Louisiana,
Lafayette
Rosemary I. Effiong, University of
Tennessee, Martin
Bryan Enderle, University of California,
Davis
David K. Erwin, Rose-Hulman Institute of
Technology
Emmanuel Ewane, Houston Community
College
Kenneth A. French, Blinn College
Donna G. Friedman, St. Louis Community
College, Florissant Valley

Herb Fynewever, Western Michigan
University
Judy George, Grossmont College
Dixie J. Goss, Hunter College City
University of New York
Ryan H. Groeneman, Jefferson College
Kimberly Hamilton-Wims, Northwest
Mississippi Community College
David Hanson, Stony Brook University
Eric Hardegree, Abilene Christian University
Michael A. Hauser, St. Louis Community
College, Meramec
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Pennsylvania
Monte L. Helm, Fort Lewis College
Sherell Hickman, Brevard Community
College
Jeffrey Hugdahl, Mercer University
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University
Richard Jarman, College of DuPage
Carolyn Sweeney Judd, Houston Community
College
Bryan King, Wytheville Community College
Peter J. Krieger, Palm Beach Community
College
John T. Landrum, Florida International
University, Miami
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University
Richard Lavallee, Santa Monica College
Debbie Leedy, Glendale Community College
Alan Levine, University of Louisiana,
Lafayette
Chunmei Li, Stephen F. Austin State University
Alan F. Lindmark, Indiana University
Northwest
Donald Linn, Indiana University Purdue
University Fort Wayne
Arthur Low, Tarleton State University
David Lygre, Central Washington University
Toni G. McCall, Angelina College
Debbie McClinton, Brevard Community
College
William McHarris, Michigan State University
Curtis McLendon, Saddleback College
Lauren McMills, Ohio University
Jennifer E. Mihalick, University of
Wisconsin, Oshkosh
John T. Moore, Stephen F. Austin State
University
Brian Moulton, Brown University
Michael R. Mueller, Rose-Hulman Institute
of Technology

My friends that make up the superb publishing team
at McGraw-Hill Higher Education have again done an
excellent job developing and producing this text. My
warmest thanks for their hard work, thoughtful advice,
and support go to Publisher Ryan Blankenship and
Executive Editor Jeff Huettman. I lost one wonderful Senior Developmental Editor, Donna Nemmers, early in the
project and found another wonderful one, Lora Neyens.
Once again, Lead Project Manager Peggy Selle created a
superb product, this time based on the clean, modern
look of Senior Designer David Hash. Marketing Manager
Tami Hodge ably presented the final text to the sales staff
and academic ­community.

siL02699_fm_i_xxvii.indd 20

Kathy Nabona, Austin Community College
Chip Nataro, Lafayette College
David S. Newman, Bowling Green State
University
William J. Nixon, St. Petersburg College
Eileen Pérez, Hillsborough Community
College
Richard Perkins, University of Louisiana,
Lafayette
Eric O. Potma, University of California,
Irvine
Nichole L. Powell, Tuskegee University
Parris F. Powers, Volunteer State Community
College
Mary C. Roslonowski, Brevard Community
College
E. Alan Sadurski, Ohio Northern University
G. Alan Schick, Eastern Kentucky University
Linda D. Schultz, Tarleton State University
Mary Sisak, Slippery Rock University
Joseph Sneddon, McNeese State University
Michael S. Sommer, University of Wyoming
Ana Maria Soto, The College of New Jersey
John E. Straub, Boston University
Richard E. Sykora, University of South
Alabama
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Stevens Point
Maria E. Tarafa, Miami Dade College
Kurt Teets, Okaloosa Walton College
Jeffrey S. Temple, Southeastern Louisiana
University
Lydia T. Tien, Monroe Community College
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Ramaiyer Venkatraman, Jackson State
University
Marie Villarba, Glendale Community College
Kirk W. Voska, Rogers State University
Edward A. Walters, University of New
Mexico
Kristine Wammer, University of St. Thomas
Shuhsien Wang-Batamo, Houston
Community College
Thomas Webb, Auburn University
Kurt Winkelmann, Florida Institute of
Technology
Steven G. Wood, Brigham Young University
Louise V. Wrensford, Albany State University
James A. Zimmerman, Missouri State
University
Susan Moyer Zirpoli, Slippery Rock University
Tatiana M. Zuvich, Brevard Community
College

Expert freelancers made indispensable contributions
as well. My superb copyeditor, Jane Hoover, continued to
improve the accuracy and clarity of my writing, and proofreaders Janelle Pregler and Angie Ruden gave their consistent polish to the final manuscript. And Jerry Marshall
helped me find the best photos, and Gary Hunt helped me
create an exciting cover.
As always, my wife Ruth was involved every step of
the way, from helping with early style decisions to checking and correcting content and layout in page proofs. And
my son Daniel consulted on the choice of photos and the
cover.

11/29/11 10:25 AM

xxi

Preface

13.3 • Solubility as an Equilibrium Process

A Guide to Student Success: How to Get the Most Out of Your Textbook
P1

7

Organizing and Focusing
Chapter Outline
The chapter begins with an outline that shows the
sequence of topics and subtopics.

A

P2

P2

Quantum Theory
and Atomic Structure
B

C

Figure 13.10 The effect of pressure on gas solubility.

Key Principles

to focus on while studying this chapter

Key Principles

a piston-cylinder assembly with a gas above a saturated aqueous solution of the gas
(Figure 13.10A). At equilibrium, at a given pressure, the same number of gas molecules enter and leave the solution per unit time:

• In a vacuum, electromagnetic radiation travels at the speed of light (c) in waves.
The properties of a wave are its wavelength (l, distance between corresponding
points on adjacent waves), frequency (n, number of cycles the wave undergoes
per second), and amplitude (the height of the wave), which is related to the
intensity (brightness) of the radiation. Any region of the electromagnetic
spectrum includes a range of wavelengths. (Section
Gas 7.1)
1 solvent BA saturated solution
• In everyday experience, energy is diffuse and matter is chunky, but certain
phenomena—blackbody radiation (the light emitted by hot objects), the
photoelectric effect (the flow of current when light strikes a metal), and atomic
spectra (the specific colors emitted from a substance that is excited)—can only be
Light from Excited Atoms In a fireworks display and
explained if energy consists of “packets” (quanta) that occur in, and thus change
by, fixed amounts. The energy of a quantum is related to its frequency.
many other everyday phenomena, we see the result of
(Section 7.1)
atoms absorbing energy and then emitting it as light. In
this chapter, we explore the basis of these phenomena
• According to the Bohr model, an atomic spectrum consists of separate lines
and learn some surprising things about the makeup of
because an atom has certain energy levels (states) that correspond to electrons
in orbits around the nucleus. The energy of the atom changes when the electron
the universe.
moves from one orbit to another as the atom absorbs (or emits) light of a specific
frequency. (Section 7.2)
gas
• Wave-particle duality means that matter has wavelike properties (as shown by
the de Broglie wavelengthgas
and electron diffraction) and energy has particle7.1
The Nature of Light
like properties (as shown by photons of light having momentum). These
Wave Nature of Light
5wekH 3 Pgas Particle Nature of Light
properties are observable only on the atomic scale, and because S
ofgas
them,
can never simultaneously know the position and speed of an electron in an atom
7.2
Atomic Spectra
(uncertainty principle). (Section 7.3)
Line Spectra and the Rydberg Equation
H
Bohr Model of the Hydrogen Atom
• According to the quantum-mechanical model of the H atom, each energy level
Energy
of the atom is associated with an atomic orbital (wave function), a mathematical
gas
gas Levels of the Hydrogen Atom
H
Spectral
Analysis
description of the electron’s position in three dimensions.
We21
can know the
21
7.3
The Wave-Particle Duality of Matter and Energy
probability that the electron is within a particular tiny volume of space, but not its
Wave Nature of Electrons and Particle Nature of Photons
exact location. The probability is highest for the electron being near the nucleus,
Heisenberg’s
Uncertainty
Principle
and it decreases with distance. (Section 7.4)
7.4
The Quantum-Mechanical Model of the Atom
• Quantum numbers denote each atomic orbital’s energy (n, principal), shape
Atomic Orbital and Probable Location of the Electron
(l, angular momentum), and spatial orientation (ml, magnetic). An energy level
Quantum Numbers of an Orbital
consists of sublevels, which consist of orbitals. There is a hierarchy of quantum
Quantum Numbers and Energy Levels
numbers: n limits l, which limits ml. (Section 7.4)
Shapes of Atomic Orbitals
• InProblem
the H atom, there
is only
one type pressure
of electrostatic of
interaction:
the attraction
The
partial
carbon
dioxide gas inside
a bottle
The Special
Case of of
the Hcola
Atom is 4 atm at
between nucleus and electron. Thus, for the H atom only, the energy levels
258C.
What
is thequantum
solubility
CO7.4)
depend
solely
on the principal
number (n).of
(Section
2? The Henry’s law constant for CO2 in water is

The main principles from the chapter are given in concise, separate paragraphs so you can keep them in mind
as you study. You may also want to review them when
you are ­finished.
9.1 • Atomic Properties and Chemical Bonds

W

hy do substances behave as they do? That is, why is table salt (or any other
ionic substance) a hard, brittle, high-melting solid that conducts a current only
when molten or dissolved in water? Why is candle wax (along with most covalent
substances) low melting, soft, and nonconducting, even though diamond (as well as
a few other exceptions) is high melting and extremely hard? And why is copper (and
most other metals) shiny, malleable, and able to conduct a current whether molten
or solid? The answers lie in the type of bonding within the substance. In Chapter 8,
we examined the properties of individual atoms and ions. But the behavior of matter
really depends on how those atoms and ions bond.

Push down on the piston, and you disturb the equilibrium: gas volume decreases, so
gas pressure (and concentration) increases, and gas particles collide with the liquid
surface more often. Thus, more particles enter than leave the solution per unit time
(Figure 13.10B). More gas dissolves to reduce this disturbance (a shift to the right in
the preceding equation) until the system re-establishes equilibrium (Figure 13.10C).
Henry’s law expresses the quantitative relationship between gas pressure and
solubility: the solubility of a gas (S ) is directly proportional to the partial pressure
Outline
of the gas (P ) above the solution:

277

(13.3)

CONCEPTS & SKILLS TO REVIEW

where k is the Henry’s law constant and is specific for a given gas-solvent combination at a given temperature. With S in mol/L and P in atm, the units of k are
mol/Latm (that is, molL atm ).

before studying this chapter

• characteristics of ionic and covalent compounds; Coulomb’s law
(Section 2.7)
• polar covalent bonds and the
polarity of water (Section 4.1)
• Hess’s law, DH 8rxn, and DH 8f
(Sections 6.5 and 6.6)
• atomic and ionic electron configurations (Sections 8.2 and 8.4)
• trends in atomic properties and
metallic behavior (Sections 8.3
and 8.4)

9.1 • Atomic ProPerties And chemicAl Bonds
Before we examine the types of chemical bonding, we should start with the most
fundamental question: why do atoms bond at all? In general, bonding lowers the
potential energy between positive and negative particles (see Figure 1.3), whether
they are oppositely charged ions or nuclei and electron pairs. Just as the strength of
attractions and repulsions among nucleus and electrons determines the properties of an
atom, the type and strength of chemical bonds determine the properties of a substance.

Sample Problem 13.2

Concepts and Skills to Review
3.331022 mol/Latm at 258C.

This
prepare
forsothe
upcoming
Planunique
We know P feature
(4 atm) and helps
the value ofyou
k (3.3310
mol/Latm),
we substitute
them into Equation 13.3 to find S .
chapter
by
referring
to
key
material
from
earlier
chapSolution S 5 k 3 P 5 (3.3310 mol/Latm)(4 atm) 5 0.1 mol/L
The you
units areshould
correct. We rounded
to one significant
figure to you
match the
numberreading
tersCheck
that
understand
before
start
in the pressure. A 0.5-L bottle of cola has about 2 g (0.05 mol) of dissolved CO .
theFollow-UP
current
one.
Problem
13.2 If air contains 78% N by volume, what is the solubility of N
CO2

CO2

in water at 258C and 1 atm (kH for N2 in H2O at 258C 5

7A 8A
(17) (18)

3A 4A 5A 6A
(13) (14) (15) (16)

Nonmetals
Metalloids

Sample Problems
Li

Be

Na

Mg

3B
(3)

4B
(4)

5B
(5)

6B
(6)

7B
(7)

H

He

•

•

B

C

N

O

F

Ne

(8)

1B 2B
8B
(9) (10) (11) (12)

Al

Si

P

S

Cl

Ar

K

Ca

Sc

Ti

V

Cr

Mn

Fe

Co

Ni

Cu

Zn

Ga

Ge

As

Se

Br

Kr

Rb

Sr

Y

Zr

Nb

Mo

Tc

Ru

Rh

Pd

Ag

Cd

In

Sn

Sb

Te

I

Xe

Cs

Ba

La

Hf

Ta

W

Re

Os

Ir

Pt

Au

Hg

Tl

Pb

Bi

Po

At

Rn

A, Location within the
periodic table. B, Relative magnitudes of
some atomic properties across a period.

A worked-out problem appears whenever an important new
concept or skill is introduced. The step-by-step approach is
shown consistently for every sample problem in the text.
• Plan analyzes the problem so that you can use what is
known to find what is unknown. This approach develops
the habit of thinking through the solution before performing
calculations.
• In many cases, a Road Map specific to the problem is shown
alongside the plan to lead you visually through the needed
calculation steps.
• Solution shows the calculation steps in the same order as
they are discussed in the plan and shown in the road map.
• Check fosters the habit of going over your work quickly
to make sure that the answer is reasonable, chemically
and mathematically—a great way to avoid careless errors.
• Comment, shown in many problems, provides an additional insight, and alternative approach, or a common mistake to avoid.
• Follow-up Problem gives you immediate practice by presenting a similar problem that requires the same approach.
Fr

Ra

Ac

Rf

Db

Sg

Bh

Hs

Mt

Ds

Rg

Cn 113 114 115 116

Ce

Pr

Nd

Pm Sm

Eu

Gd

Tb

Dy

Ho

Er

Tm

Yb

Lu

Th

Pa

U

Np

Am Cm

Bk

Cf

Es

Fm

Md

No

Lr

A

siL02699_fm_i_xxvii.indd 21

Pu

118

B

PROPERTY

METAL
ATOM

NONMETAL
ATOM

Atomic size

Larger

Smaller

Zeff

Lower

Higher

IE

Lower

Higher

EA

Less negative More negative

mol/Latm)?

2

Summary of Section 13.3

Step-By-Step Problem Solving
Using this clear and thorough problem-solving approach,
you’ll learn to think through chemistry problems logically and
Figure 9.1 A comparison of metals
­systematically.
and nonmetals.

731024

216

dencies to lose or gain electrons. Such differences occur between reactive metals
[Groups 1A(1) and 2A(2)] and nonmetals [Group 7A(17) and the top of Group 6A(16)].
A metal atom (low IE) loses its one or two valence electrons, and a nonmetal atom
(highly negative EA) gains the electron(s). Electron transfer from metal to nonmetal
occurs, and each atom forms an ion with a noble gas electron configuration. The electrostatic attractions between these positive and negative ions draw them into a threedimensional array to form an ionic solid. Note that the chemical formula of an ionic
compound is the empirical formula because it gives the cation-to-anion ratio.

Metals

CO2

2

A bulleted list of statements conclude each section, immedi1. Metal
with nonmetal: electron
and ionic
bonding (Figure
next
ately
reiterating
thetransfer
major
ideas
just9.2A,
covered.
page). We observe ionic bonding between atoms with large differences in their ten-

2A
(2)

H

22

2

In general, there is a gradation from atoms of more metallic elements to atoms of
more nonmetallic elements across a period and up a group (Figure 9.1). Three types
of bonding result from the three ways these two types of atoms can combine:

Key:

22

H

CO2

Section
Summaries
Types of Bonding:
Three Ways Metals and Nonmetals Combine

1A
(1)

Using Henry’s Law to Calculate Gas Solubility

 solutionthatcontainsthemaximumamountofdissolvedsoluteinthepresence
A
of excess undissolved solute is saturated. A saturated solution is in equilibrium
with excess solute, because solute particles are entering and leaving the solution
at the same rate.
Most solids are more soluble at higher temperatures.

•

AllgaseshaveanegativeDHsolninwater,soheatinglowersgassolubilityinwater.

•

 enry’s law says that the solubility of a gas is directly proportional to its partial
H
pressure above the solution.
3.5 • Fundamentals of Solution Stoichiometry

103

Stoichiometry of Reactions in Solution
Solving stoichiometry problems for reactions in solution requires the additional step
of converting the volume of reactant or product in solution to amount (mol):
1.
2.
3.
4.

Balance the equation.
Find the amount (mol) of one substance from the volume and molarity.
Relate it to the stoichiometrically equivalent amount of another substance.
Convert to the desired units.

Sample Problem 3.25

 Calculating Quantities of Reactants and Products

for a Reaction in Solution

Problem Specialized cells in the stomach release HCl to aid digestion. If they release
too much, the excess can be neutralized with an antacid. A common antacid contains
magnesium hydroxide, which reacts with the acid to form water and magnesium chloride
solution. As a government chemist testing commercial antacids, you use 0.10 M HCl to
simulate the acid concentration in the stomach. How many liters of “stomach acid” react
with a tablet containing 0.10 g of magnesium hydroxide?
Plan We are given the mass (0.10 g) of magnesium hydroxide, Mg(OH)2, that reacts with
the acid. We also know the acid concentration (0.10 M) and must find the acid volume.
After writing the balanced equation, we convert the mass (g) of Mg(OH)2 to amount (mol)
and use the molar ratio to find the amount (mol) of HCl that reacts with it. Then, we use
the molarity of HCl to find the volume (L) that contains this amount (see the road map).
Solution Writing the balanced equation:
Mg(OH)2(s) 1 2HCl(aq) -£ MgCl2(aq) 1 2H2O(l)
Converting from mass (g) of Mg(OH)2 to amount (mol):
1 mol Mg(OH)2
Amount (mol) of Mg(OH)2 5 0.10 g Mg(OH)2 3
5 1.731023 mol Mg(OH)2
58.33 g Mg(OH)2
Converting from amount (mol) of Mg(OH)2 to amount (mol) of HCl:
2 mol HCl
5 3.431023 mol HCl
1 mol Mg(OH)2
Converting from amount (mol) of HCl to volume (L):
1L
5 3.431022 L
Volume (L) of HCl 5 3.431023 mol HCl 3
0.10 mol HCl
Check The size of the answer seems reasonable: a small volume of dilute acid (0.034 L
of 0.10 M) reacts with a small amount of antacid (0.0017 mol).
Comment In Chapter 4, you’ll see that this equation is an oversimplification, because
HCl and MgCl2 exist in solution as separated ions.

Amount (mol) of HCl 5 1.731023 mol Mg(OH)2 3

Road Map
Mass (g) of Mg(OH)2
divide by  (g/mol)

Amount (mol) of Mg(OH)2
molar ratio

Amount (mol) of HCl
divide by M (mol/L)

Volume (L) of HCl

Follow-UP Problem 3.25 Another active ingredient in some antacids is aluminum
hydroxide. Which is more effective at neutralizing stomach acid, magnesium hydroxide or
aluminum hydroxide? [Hint: “Effectiveness” refers to the amount of acid that reacts with
a given mass of antacid. You already know the effectiveness of 0.10 g of Mg(OH)2.]

Except for the additional step of finding amounts (mol) in solution, limitingreactant problems for reactions in solution are handled just like other such problems.

Sample Problem 3.26

 Solving Limiting-Reactant Problems for Reactions

in Solution

Problem Mercury and its compounds have uses from fillings for teeth (as a mixture with
silver, copper, and tin) to the production of chlorine. Because of their toxicity, however,
soluble mercury compounds, such as mercury(II) nitrate, must be removed from industrial
wastewater. One removal method reacts the wastewater with sodium sulfide solution to
produce solid mercury(II) sulfide and sodium nitrate solution. In a laboratory simulation,
0.050 L of 0.010 M mercury(II) nitrate reacts with 0.020 L of 0.10 M sodium sulfide.
(a) How many grams of mercury(II) sulfide form? (b) Write a reaction table for this process.

11/29/11 10:25 AM

403

percent) of dissociated HA molecules that increases with dilution.

NK OF IT THIS WAY

Weak acids dissociating to a greater extent as they are diluted is analogous to the
shift in a gaseous reaction when the container volume increases (Section 17.6). In
the gaseous reaction, an increase in volume as the piston is withdrawn shifts the
equilibrium position to favor more moles of gas. In the case of HA dissociation, the
increase in volume as solvent is added shifts the equilibrium position to favor more
moles of ions.

Are Gaseous and
-Acid Equilibria Alike?

xxii

Preface

Sample Problem 18.9 uses molecular scenes to highlight this idea. (Note that, in
order to depict the scenes practically, the acid has a much higher percent dissociation
than any real weak acid does.)

Sample Problem 18.9

UsingMolecularScenestoDeterminetheExtent
ofHADissociation

Problem A 0.15 M solution of HA (blue and green) is 33% dissociated. Which scene
represents a sample of that solution after it is diluted with water?

Chapter 16 • Chapter Review Guide

Key Terms

535

These important terms appear in boldface in the chapter and are defined again in the Glossary.

Section 16.1
chemical kinetics (499)
reaction rate (499)
Section 16.2
average rate (502)
instantaneous rate (503)
initial rate (503)
Section 16.3
rate law (rate equation) (505)
rate constant (505)
reaction orders (505)

Section 16.4
integrated rate law (512)
half-life (t1/2) (515)
Section 16.5
collision theory (519)
Arrhenius equation (519)
activation energy (Ea) (520)
effective collision (522)
frequency factor (522)
transition state theory (522)
transition state (activated
complex) (522)
reaction energy diagram
(523)

Section 16.6
reaction mechanism (525)
elementary reaction
(elementary step) (526)
molecularity (526)
unimolecular reaction (526)
bimolecular reaction (526)
rate-determining (ratelimiting) step (527)
reaction intermediate (527)

Section 16.7
catalyst (530)
homogeneous catalyst (531)
heterogeneous catalyst (532)
hydrogenation (532)
enzyme (533)
active site (533)

Unique to Principles of General Chemistry:
Molecular-Scene Sample Problems

These problems apply the same stepwise strategy to help
you interpret molecular scenes and solve problems based
Key Equations and Relationships
on them.
16.1 Expressing reaction rate in terms of reactant A (501):
16.6 Calculating the time to reach a given [A] in a zero-order
Numbered and screened concepts are listed for you to refer to or memorize.

Rate 5 2
1

2

3

Plan We are given the percent dissociation of the original HA solution (33%), and we
know that the percent dissociation increases as the acid is diluted. Thus, we calculate
the percent dissociation of each diluted sample and see which is greater than 33%.
To determine percent dissociation, we apply Equation 18.5, with HAdissoc equal to the
number of H3O1 (or A2) and HAinit equal to the number of HA plus the number of
H3O1 (or A2).
Solution Calculating the percent dissociation of each diluted solution with Equation 18.5:
Solution 1. Percent dissociated 5 4/(5 1 4) 3 100 5 44%
Solution 2. Percent dissociated 5 2/(7 1 2) 3 100 5 22%
Solution 3. Percent dissociated 5 3/(6 1 3) 3 100 5 33%

reaction (rate 5 k) (513):

D 3A 4
Dt

3A 4 t 2 3A 4 0 5 2kt
16.7 Finding the half-life of a first-order process (516):
ln 2
0.693
t1/2 5
5
k
k
16.8 Relating the rate constant to the temperature (Arrhenius
equation) (519):

16.2 Expressing the rate of a general reaction (503):
1 D 3B 4
1 D 3C 4
1 D 3D 4
1 D 3A 4
52
5
5
Rate 5 2
c Dt
a Dt
b Dt
d Dt
16.3 Writing a general rate law (in which products do not appear)
(505):
Rate 5 k 3A 4 m 3B 4 n . . .

k 5 Ae2Ea /RT
16.9 Relating the heat of reaction to the forward and reverse activation energies (520):
DHrxn 5 Ea(fwd) 2 Ea(rev)

16.4 Calculating the time to reach a given [A] in a first-order reaction (rate 5 k[A]) (513):
3A 4 0
ln
5 kt
3A 4

16.10 Calculating the activation energy (rearranged form of
Arrhenius equation) (521):
Ea 1
k2
1
ln 5 2 a 2 b
k1
R T2
T1

t

16.5 Calculating the time to reach a given [A] in a simple secondorder reaction (rate 5 k[A]2) (513):

Therefore, scene 1 represents the diluted solution.
Check Let’s confirm our choice by examining the other scenes: in scene 2, HA is
less dissociated than originally, so that scene must represent a more concentrated HA
solution; scene 3 represents another solution with the same percent dissociation as the
original.

FOllOW-UP PrOblem 18.9 The scene in the margin represents a sample of a weak acid
HB (blue and purple) dissolved in water. Draw a scene that represents the same volume
after the solution has been diluted with water.

1
1
2
5 kt
3A 4 t
3A 4 0

Brief SolutionS to follow-up proBlemS
16.1

Compare your own solutions to these calculation steps and answers.

(a) 4NO(g) 1 O2(g) -£ 2N2O3(g);

rate 5 2

D 3O2 4
1 D 3NO 4
1 D 3N2O3 4
52
5
Dt
4 Dt
2
Dt

D 3O2 4
1 D 3NO4
1
(b) 2
52
5 2 (21.60310 mol/Ls)
Dt
4 Dt
4
5 4.00310 25 mol/Ls

16.2

Brief Solutions to Follow-up Problems

First order in Br2, first order in BrO32, second order in H1,
fourth order overall.

These provide multistep solutions at the end of the
chapter, not just a one-number answer at the back of
the book. This fuller treatment provides an excellent
way for you to reinforce problem-solving skills.

16.4 (a) The rate law shows the reaction is zero order in Y,
so the rate is not affected by doubling Y: rate of Expt 2 5 0.25310–5
mol/Ls.
(b) The rate of Expt 3 is four times that of Expt 1, so [X]
doubles.

16.3

Rate 5 k[H2]m[I2]n. From Expts 1 and 3, m 5 1. From Expts 2
and 4, n 5 1. Therefore, rate 5 k[H2][I2]; second order overall.

4.6 • Elements in Redox Reactions

16.5 1/[HI]1 2 1/[HI]0 5 kt
111 L/mol 2 100. L/mol 5 (2.4310221 L/mols)(t)
t 5 4.631021 s (or 1.531014 yr)
16.6 (a)

139

Visualizing Chemistry
Figure 4.15 A more reactive metal (Cu)
displacing the ion of a less reactive metal
(Ag) from solution.

Three-Level Illustrations
A Silberberg hallmark, these illustrations provide macroscopic
and molecular views of a process to help you connect these two
levels of reality with each other and with the chemical equation
that describes the process in symbols.

Copper wire
coated with
silver

Copper
wire

Copper
nitrate
solution

Silver
nitrate
solution

12.6 • The Solid State: Structure, Properties, and Bonding

375

Cu2 +
Ag+

A Simple cubic

B Body-centered cubic

C Face-centered cubic

Ag+

2e–
Ag atoms
coating
wire

Cu atoms
in wire
0
+1+5 –2
2AgNO3(aq) + Cu(s)

5 –2
0
+2 +5
–2
Cu(NO3)2(aq) + 2Ag(s)

Consider the metals we’ve just discussed: Li, Al, and Ni lie above H2, while Ag
lies below it; also, Zn lies above Cu, which lies above Ag. The most reactive metals
on the list are in Groups 1A(1) and 2A(2) of the periodic table, and the least reactive
are Cu, Ag, and Au in Group 1B(11) and Hg in 2B(12).

Cutting-Edge Molecular Models
2

2

2

2

Thus, elemental chlorine can oxidize bromide ions (below) or iodide ions from solution, and elemental bromine can oxidize iodide ions:
1

0

0

1

2Br(aq)  Cl2(aq) ±£ Br2(aq)  2Cl(aq)

Combustion Reactions
​ ombustion is the process of combining with oxygen, most commonly with the release
C

of heat and the production of light, as in a flame. Combustion reactions are not
classified by the number of reactants and products, but all of these reactions are redox
processes because elemental oxygen is a reactant:
2CO(g) 1 O2(g) -£ 2CO2(g)
The combustion reactions that we commonly use to produce energy involve coal,
petroleum, gasoline, natural gas, or wood as a reactant. These mixtures consist of
substances with many C-C and C-H bonds, which break during the reaction, and
siL02699_fm_i_xxvii.indd 22
each C and H atom combines with oxygen to form CO and H O. The combustion

Strength as reducing agent

Author
andSeries
artist
worked
byof side
and decreases
employed
The Activity
of the
Halogens side
Reactivity
the elements
down the
Group 7A(17), so we have
most
advanced computer-graphic software to provide accuF . Cl . Br . I
rate
molecular-scale models and vivid scenes.
A halogen higher in the group is a stronger oxidizing agent than one lower down.

Li
K
Can displace H2
Ba
from water
Ca
Na
Mg
Al
Mn
Coordination number = 8
Coordination number = 12
Coordination
number = 6
Can displace
H2
Zn
from steam
1
1
1
–
–
–
8 atom
Cr
8 atom
8 atom
at
8
corners
at 8 corners
at 8 corners
Fe
Cd
1
–
1 atom
Co
2 atom
at 6 faces
at center
Ni
Can displace H2
from acid
Sn
1
1
1
1
Atoms/unit cell = ( 8– x 8) + 1 = 2
Atoms/unit cell = ( 8– x 8) + ( 2– x 6) = 4
Atoms/unit cell = 8– x 8 = 1
Pb
H2
Figure 12.23 The three cubic unit cells. A, Simple cubic unit cell. face-centered atoms yellow. Third row: A unit cell (shaded blue) in an
B, Body-centered cubic unit cell. C, Face-centered cubic unit cell. Top
expanded portion of the crystal. The number of nearest neighbors around
Cu
row: Cubic arrangements of atoms in expanded view. Second row:
one particle (dark blue in center) is the coordination number. Bottom row:
Hg Cannot
Space-filling
view of
cubic arrangements. All atoms are identical
The total numbers of atoms in the actual unit cell. The simple cubic has
displace
Hthese
2
but, for
clarity,
corner atoms are blue, body-centered atoms pink, and
one atom, the body-centered has two, and the face-centered has four.
from
any
source
Ag
Au

Figure 4.16 The activity series of the

metals. The most active metal (strongest
reducing agent) is at the top, and the least
active metal (weakest reducing agent) is
at the bottom.

11/29/11 10:25 AM

Preface

xxiii
417

Chapter 13 • Chapter Review Guide

Reinforcing The Learning 
Process

chapter review Guide
Learning Objectives

Chapter Review Guide

1. Explainhowsolubilitydependsonthetypesofintermolecularforces(like-dissolves-likerule)andunderstandthecharacteristicsofsolutionsconsistingofgases,liquids,orsolids
(§13.1)(SP13.1)(EPs13.1–13.12)
2. UnderstandtheenthalpycomponentsofDHsoln,the
dependenceofDHhydronchargedensity,andwhyasolution
processisexothermicorendothermic(§13.2)(EPs13.13–
13.15,13.18–13.25,13.28)
3. Comprehendthemeaningofentropyandhowthebalance
betweenDHandDSgovernsthesolutionprocess(§13.2)
(EPs13.16,13.17,13.26,13.27)
4. Distinguishamongsaturated,unsaturated,andsupersaturated
solutionsandexplaintheequilibriumnatureofasaturated
solution(§13.3)(EPs13.29,13.35)

• Learning Objectives are listed, with section, sample problem,
and end-of-chapter problem numbers, to help you focus on key
concepts and skills.
• Key Terms are boldfaced within the chapter and listed here by
section (with page numbers); they are defined again in the
Glossary.
• Key Equations and Relationships are highlighted and numbered within the chapter and listed here with page numbers.

Key Terms

Chapter 13 • The Properties of Solutions

13.79 What is the minimum mass of glycerol (C3H8O3) that must
be dissolved in 11.0 mg of water to prevent the solution from
freezing at 2158C? (Assume ideal behavior.)
180 Calculate the molality and van’t Hoff factor (i) for the following aqueous solutions:
(a) 1.00 mass % NaCl, freezing point 5 20.5938C
(b) 0.500 mass % CH3COOH, freezing point 5 20.1598C
13.81 Calculate the molality and van’t Hoff factor (i) for the following aqueous solutions:
(a) 0.500 mass % KCl, freezing point 5 20.2348C
(b) 1.00 mass % H2SO4, freezing point 5 20.4238C
13.82 In a study designed to prepare new gasoline-resistant coatings, a polymer chemist dissolves 6.053 g of poly(vinyl alcohol) in
enough water to make 100.0 mL of solution. At 258C, the osmotic
pressure of this solution is 0.272 atm. What is the molar mass of
the polymer sample?
13.83 The U.S. Food and Drug Administration lists dichloromethane (CH2Cl2) and carbon tetrachloride (CCl4) among the
many cancer-causing chlorinated organic compounds. What are
the partial pressures of these substances in the vapor above a solution of 1.60 mol of CH2Cl2 and 1.10 mol of CCl4 at 23.58C? The
vapor pressures of pure CH2Cl2 and CCl4 at 23.58C are 352 torr
and 118 torr, respectively. (Assume ideal behavior.)

Comprehensive Problems
13.84 The three aqueous ionic solutions represented below have
total volumes of 25. mL for A, 50. mL for B, and 100. mL for
C. If each sphere represents 0.010 mol of ions, calculate: (a) the
total molarity of ions for each solution; (b) the highest molarity
of solute; (c) the lowest molality of solute (assuming the solution
densities are equal); (d) the highest osmotic pressure (assuming
ideal behavior).
+

–
–

+
–

–
+

+
+

–

–

+
+

2+

–

–

2+

–

–
2+

–

+

B

A

–

–

–

–

+

–

C

2+

–

13.85 Gold occurs in seawater at an average concentration of
1.131022 ppb. How many liters of seawater must be processed to
recover 1 troy ounce of gold, assuming 81.5% efficiency (d of seawater 5 1.025 g/mL; 1 troy ounce 5 31.1 g)?
13.86 Use atomic properties to explain why xenon is 11 times as
soluble as helium in water at 08C on a mole basis.
13.87 Which of the following best represents a molecular-scale
view of an ionic compound in aqueous solution? Explain.
+

+

+

+

A

+
–
–
+
B

–

+

+

–

These are concepts and skills to review after studying this chapter.

Related section (§), sample problem (SP), and upcoming end-of-chapter
problem (EP) numbers are listed in parentheses.

A rich catalog of study aids ends each chapter to help
you review its content:

422

The following sections provide many aids to help you study this chapter. (Numbers in
parentheses refer to pages, unless noted otherwise.)

C

13.88 Four 0.50 m aqueous solutions are depicted. Assume the
solutions behave ideally: (a) Which has the highest boiling point?

These important terms appear in boldface in the chapter and are defined again in the Glossary.

Section 13.1
solute(392)
solvent(392)
miscible(392)
solubility(S)(392)
like-dissolves-likerule(393)
hydrationshell(393)
(b) Which has the lowest freezing point? (c) Can you determine
ion–induceddipole
which one has the highest osmotic pressure? Explain.
force(393)
2+
dipole–induceddipole
–
–
+
–
+
force(393)
2+
–
–
–
–
+
alloy(396)
–
2+
–
–
–
Section
13.2
+
2+
–
A
B
heatofsolution
(DH
soln)(397)
2–

solvation(397)
hydration(398)
heatofhydration
(DHhydr)(398)
chargedensity(398)
entropy(S)(399)
Section 13.3
saturatedsolution(401)
unsaturatedsolution(401)
supersaturatedsolution(401)
Henry’slaw(403)
Section 13.4
molality(m)(404)
masspercent[%(w/w)]
(405)

boilingpointelevation
(DTb)(410)
freezingpointdepression
(DTf)(411)
semipermeablemembrane
(412)
osmosis(412)
osmoticpressure
(P)(413)
ionicatmosphere(415)

2–
2+

Key Equations and Relationships

2+
2–

C

Numbered and screened concepts are listed for you to refer to or memorize.

D

13.1 Dividingthegeneralheatofsolutionintocomponent
13.89 “De-icing salt” is used to melt snow and ice on streets. The enthalpies(397):
highway department of a small town is deciding whether to buy
NaCl or CaCl2, which are equally effective, to use for this purpose.
The town can obtain NaCl for $0.22/kg. What is the maximum the
town should pay for CaCl2 to be cost effective?

13.90 Thermal pollution from industrial wastewater causes the temperature of river or lake water to increase, which can affect fish
survival as the concentration of dissolved O2 decreases. Use the following data to find the molarity of O2 at each temperature (assume
the solution density is the same as water):
Temperature
(°C)

Solubility of O2
(mg/kg H2O)

0.0
20.0
40.0

Density of
H2O (g/mL)

14.5
9.07
6.44

0.99987
0.99823
0.99224

13.91 A chemist is studying small organic compounds for their
potential use as an antifreeze. When 0.243 g of a compound is
dissolved in 25.0 mL of water, the freezing point of the solution is 20.2018C. (a) Calculate the molar mass of the compound
(d of water 5 1.00 g/mL). (b) Analysis shows that the compound
is 53.31 mass % C and 11.18 mass % H, the remainder being O.
Calculate the empirical and molecular formulas of the compound.
(c) Draw a Lewis structure for a compound with this formula that
forms H bonds and another for one that does not.
13.92 Is 50% by mass of methanol dissolved in ethanol different
from 50% by mass of ethanol dissolved in methanol? Explain.
13.93 Three gaseous mixtures of N2 (blue), Cl2 (green), and Ne
(purple) are depicted below. (a) Which has the smallest mole fraction of N2? (b) Which have the same mole fraction of Ne? (c) Rank
all three in order of increasing mole fraction of Cl2.

A

volumepercent
[%(v/v)](405)
molefraction(X)(405)
Section 13.5
colligativeproperty(408)
electrolyte(408)
nonelectrolyte(408)
vaporpressurelowering
(DP)(408)
Raoult’slaw(409)
idealsolution(409)

2+

2+
2–

5. Describetheeffectoftemperatureonthesolubilityof
solidsandgasesinwaterandtheeffectofpressureonthe
solubilityofgases(Henry’slaw)(§13.3)(SP13.2)(EPs
13.30–13.34,13.36)
6. Expressconcentrationintermsofmolarity,molality,mole
fraction,andpartsbymassorbyvolumeandbeableto
interconverttheseterms(§13.4)(SPs13.3–13.5)(EPs
13.37–13.58)
7. Describeelectrolytebehaviorandthefourcolligative
properties,explainthedifferencebetweenphasediagrams
forasolutionandapuresolvent,explainvapor-pressure
loweringfornon-volatileandvolatilenonelectrolytes,and
discussthevan’tHofffactorforcolligativepropertiesof
electrolytesolutions(§13.5)(SPs13.6–13.9)(EPs13.59–
13.83)

B

C

13.94 Four U tubes each have distilled water in the right arm, a
solution in the left arm, and a semipermeable membrane between
arms. (a) If the solute is KCl, which solution is most concentrated?

13.5 Definingconcentrationintermsofmolality(404):
Molality(m) 5

DHsoln5DHsolute1DHsolvent1DHmix
13.2 Dividingtheheatofsolutionofanioniccompoundinwater
intocomponententhalpies(398):

13.6 Definingconcentrationintermsofmasspercent(405):
Masspercent3%(w/w)4 5

DHsoln5DHlattice1DHhydroftheions
13.3 Relatinggassolubilitytoitspartialpressure(Henry’slaw)
(403):
Sgas5kH3Pgas

massofsolute
3 100
massofsolution

13.7 Definingconcentrationintermsofvolumepercent(405):

13.4 Definingconcentrationintermsofmolarity(404):
Molarity(M) 5

amount(mol)ofsolute
mass(kg)ofsolvent

Volumepercent3%(v/v)4 5

volumeofsolute
3 100
volumeofsolution

amount(mol)ofsolute
volume(L)ofsolution

End-of-Chapter Problems
An exceptionally large number of problems ends each
chapter. These are sorted by section, and many are grouped
in similar pairs, with one of each pair answered in Appendix E (along with other problems having a colored number).
Following these section-based problems is a large group of
Comprehensive Problems, which are based on concepts and
skills from any section and/or earlier chapter and are filled
with applications from related sciences.
21.2 • Voltaic Cells: Using Spontaneous Reactions to Generate Electrical Energy

Here are some memory aids to help you connect the half-reaction with its electrode:
1. The words anode and oxidation start with vowels; the words cathode and re­duction­start
with consonants.
2. Alphabetically, the A in anode comes before the C in cathode, and the O in oxidation
comes before the R in reduction.
3. Look at the first syllables and use your imagination:

Think of It This Way
Analogies, memory shortcuts, and new insights
into key ideas are provided in “Think of It This
Way” features.

693

THINK OF IT THIS WAY
Which Half-Reaction Occurs
at Which Electrode?

ANode, OXidation; REDuction, CAThode =: AN OX and a RED CAT

Summary of Section 21.1
•

 n oxidation-reduction (redox) reaction involves the transfer of electrons from a
A
reducing agent to an oxidizing agent.

•

 he half-reaction method of balancing divides the overall reaction into halfT
reactions that are balanced separately and then recombined.

•

 herearetwotypesofelectrochemicalcells.Inavoltaiccell,aspontaneousreaction
T
generateselectricityanddoesworkonthesurroundings.Inanelectrolyticcell,the
surroundingssupplyelectricitythatdoesworktodriveanonspontaneousreaction.

•

Inbothtypesofcell,twoelectrodesdipintoelectrolytesolutions;oxidationoccurs
at the anode, and reduction occurs at the cathode.

21.2 • Voltaic cells: Using spontaneoUs
Reactions to geneRate electRical
eneRgy
siL02699_fm_i_xxvii.indd 23

When you put a strip of zinc metal in a solution of Cu21 ion, the blue color of the
solution fades and a brown-black crust of Cu metal forms on the Zn strip (Figure
21.3). During this spontaneous reaction, Cu21 ion is reduced to Cu metal, while Zn
metal is oxidized to Zn21 ion. The overall reaction consists of two half-reactions:

12/1/11 8:33 AM

xxiv

Preface

McGraw-Hill Connect®Chemistry

back are all saved online—so students can continually
review their progress and plot their course to success.

ConnectPlus® Chemistry
www.mcgraw-hillconnect.com/chemistry
McGraw-Hill Connect® Chemistry is a web-based assignment and assessment platform that gives students the means
to better connect with their coursework, with their instructors, and with the important concepts that they will need to
know for success now and in the future. The chemical drawing tool found within Connect Chemistry is CambridgeSoft’s ChemDraw®, which is widely considered the gold
standard of scientific drawing programs and the cornerstone
application for scientists who draw and annotate molecules,
reactions, and pathways. This collaboration of Connect and
ChemDraw features an easy-to-use, intuitive, and comprehensive course management and homework system with
professional-grade drawing capabilities.
With Connect Chemistry, instructors can deliver assignments, quizzes, and tests online. Questions from the text are
presented in an auto-gradable format and tied to the text’s
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and author entirely new problems. They also can track individual student performance—by question, assignment, or in
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By choosing Connect Chemistry, instructors are providing their students with a powerful tool for improving
academic performance and truly mastering course material.
Connect Chemistry allows students to practice important
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siL02699_fm_i_xxvii.indd 24

Like Connect Chemistry, McGraw-Hill ConnectPlus®
Chemistry provides students with online assignments and
assessments, plus 24/7 online access to an eBook—an
online edition of the text—to aid them in successfully completing their work, wherever and whenever they choose.

LearnSmart™
This adaptive diagnostic learning system, powered by Connect Chemistry and based on artificial intelligence, constantly assesses a student’s knowledge of the course
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innovative study tool also has features to allow the instructor to see exactly what students have accomplished, with a
built-in assessment tool for graded assignments. You can
access LearnSmart for Chemistry at www.mcgrawhillcon
nect.com/chemistry.

McGraw-Hill Higher Education
and Blackboard®
McGraw-Hill Higher Education and
Blackboard have teamed up! What
does this mean for you?
1. Your life, simplified. Now you and your students can
access McGraw-Hill’s Connect and Create right from
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2. Deep integration of content and tools. Not only do you
get single sign-on with Connect and Create, you also get
deep integration of McGraw-Hill content and content
engines right in Blackboard. Whether you’re choosing a
book for your course or building Connect assignments, all
the tools you need are right where you want them—inside
of Blackboard.
3. Seamless Gradebooks. Are you tired of keeping multiple
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center.

11/29/11 10:25 AM

Preface

4. A solution for everyone. Whether your institution is
already using Blackboard or you just want to try Blackboard on your own, we have a solution for you. McGrawHill and Blackboard can now offer you easy access to
industry leading technology and content, whether your
campus hosts it, or we do. Be sure to ask your local
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Customizable Textbooks: Create™
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• Photos The photo collection contains digital files of photographs from the text, which can be reproduced for multiple classroom uses.

siL02699_fm_i_xxvii.indd 25

xxv

• Tables Every table that appears in the text has been saved
in electronic form for use in classroom presentations and/
or quizzes.
• Animations Numerous full-color animations illustrating
important processes are also provided. Harness the visual
impact of concepts in motion by importing these files into
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Digital Lecture Capture: Tegrity records and distributes
your lecture with just a click of a button. Students can view
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Tegrity indexes as it records your slideshow presentations
and anything shown on your computer, so students can use
keywords to find exactly what they want to study.
Computerized Test Bank Prepared by Walter Orchard,
Professor Emeritus of Tacoma Community College, over
2300 test questions to accompany Principles of General
Chemistry are available utilizing Brownstone’s Diploma
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Instructors’ Solutions Manual This supplement, prepared by Patricia Amateis of Virginia Tech, contains complete, worked-out solutions for all the end-of-chapter
problems in the text. It can be found within the Instructors
Resources, on the Connect: Chemistry site.
Content Delivery Flexibility Principles of General
Chemistry, by Martin Silberberg, is available in other formats
in addition to the traditional textbook, giving instructors and
students more choices for the format of their chemistry text.
Cooperative Chemistry Laboratory Manual Prepared
by Melanie Cooper of Clemson University, this innovative
manual features open-ended problems designed to simulate
experience in a research lab. Working in groups, students
investigate one problem over a period of several weeks, so
they might complete three or four projects during the semester, rather than one preprogrammed experiment per class.
The emphasis is on experimental design, analytic problem
solving, and communication.

11/29/11 10:25 AM

xxvi

Preface

learning system resources
For Students
Student Study Guide This valuable study guide, prepared
by Libby Bent Weberg, is designed to help you recognize
your learning style; understand how to read, classify, and
create a plan for solving a problem; and practice your
problem-solving skills. For each section of each chapter, the
guide provides study objectives and a summary of the corresponding text. Following the summary are sample problems with detailed solutions. Each chapter has true-false
questions and a self-test, with all answers provided at the
end of the chapter.
Student Solutions Manual This supplement, prepared
by Patricia Amateis of Virginia Tech, contains detailed solutions and explanations for all problems in the main text that
have colored numbers.

Connect Chemistry With Connect Chemistry, you can
practice solving assigned homework problems using the
Silberberg problem-solving methodology applied in the

siL02699_fm_i_xxvii.indd 26

textbook. Algorithmic problems serve up multiple versions of similar problems for mastery of content, with
hints and feedback for common incorrect answers to help
you stay on track. Where appropriate, you engage in accurate, professional-grade chemical drawing through the use
of CambridgeSoft’s ChemDraw tool, which is implemented directly into homework problems. Check it out at
www.mcgrawhillconnect.com/chemistry.

LearnSmart/ This adaptive diagnostic learning system,
powered by Connect: Chemistry and based on artificial
intelligence, constantly assesses your knowledge of
the course material. As you work within the system,
Learn­Smart develops a personal learning path adapted to
what you have actively learned and retained. This innovative study tool also has features to allow your instructor
to see exactly what you have accomplished, with a builtin assessment tool for graded assignments. You can
access Learn­Smart for general chemistry by going to
www.mcgrawhillconnect.com/chemistry.
Animations for MP3/iPod A number of animations are
available for download to your MP3/iPod through the textbook’s Connect website.

11/29/11 10:25 AM

Principles of

GENERAL
CHEMISTRY

siL02699_fm_i_xxvii.indd 1

12/2/11 1:14 PM

1

Keys to the Study
of Chemistry

Key Principles
to focus on while studying this chapter
• M
 atter can undergo two kinds of change: physical change involves a change in
state—gas, liquid, or solid—but not in ultimate makeup (composition); chemical
change (reaction) is more fundamental because it does involve a change in
composition. The changes we observe result ultimately from changes too small
to observe. (Section 1.1)
• Energy occurs in different forms that are interconvertible, even as the total
quantity of energy is conserved. When opposite charges are pulled apart, their
potential energy increases; when they are released, potential energy is converted
to the kinetic energy of the charges moving together. Matter consists of charged
particles, so changes in energy accompany changes in matter. (Section 1.1)
• The scientific method is a way of thinking that involves making observations and
gathering data to develop hypotheses that are tested by controlled experiments
until enough results are obtained to create a model (theory) that explains an
aspect of nature. A sound theory can predict events but must be changed if
new results conflict with it. (Section 1.2)
• Any measured quantity is expressed by a number together with a unit.
Conversion factors are ratios of equivalent quantities having different units; they
are used in calculations to change the units of quantities. Decimal prefixes and
exponential notation are used to express very large or very small quantities.
(Section 1.3)
• The SI system consists of seven fundamental units, each identifying a physical
quantity such as length (meter), mass (kilogram), or temperature (kelvin). These
are combined into many derived units used to identify quantities such as volume,
density, and energy. Extensive properties, such as mass, depend on sample size;
intensive properties, such as temperature, do not. (Section 1.4)
• Uncertainty characterizes every measurement and is indicated by the number of
significant figures. We round the final answer of a calculation to the same number
of digits as in the least certain measurement. Accuracy refers to how close a
measurement is to the true value; precision refers to how close measurements
are to one another. (Section 1.5)

A Molecular View Within a Storm Lightning supplies
the energy for many atmospheric chemical changes to
occur. In fact, all the events within and around you have
causes and effects at the atomic level of reality.

Outline
1.1

Some Fundamental Definitions
Properties of Matter
States of Matter
Central Theme in Chemistry
Importance of Energy

1.2
1.3

The Scientific Approach: ­Developing a Model
Chemical Problem Solving
Units and Conversion Factors
A Systematic Approach

1.4

Measurement in Scientific Study
Features of SI Units
SI Units in Chemistry
Extensive and Intensive Properties

1.5

Uncertainty in Measurement: ­Significant Figures
Determining Significant Digits
Calculations and Rounding Off
Precision, Accuracy, and Instrument Calibration

2

siL02699_ch01_002_031.indd 2

9/28/11 1:44 PM

1.1 • Some Fundamental Definitions    3

aybe you’re taking this course because chemistry is fundamental to understanding
other natural sciences. Maybe it’s required for your major. Or maybe you just
want to learn more about the impact of chemistry on society or even on your everyday
life. For example, did you have cereal, fruit, and coffee for breakfast today? In chemical
terms, you enjoyed nutrient-enriched, spoilage-retarded carbohydrate flakes mixed in a
white emulsion of fats, proteins, and monosaccharides, with a piece of fertilizer-grown,
pesticide-treated fruit, and a cup of hot aqueous extract of stimulating alkaloid. Earlier,
you may have been awakened by the sound created as molecules aligned in the liquidcrystal display of your clock and electrons flowed to create a noise. You might have
thrown off a thermal insulator of manufactured polymer and jumped in the