Chemistry The Cen t ral S c ien c e 13^TH Edition

Chemistry The Cen t ral S c ien c e 13^TH Edition

Theodore L. Brown

University of Illinois at Urbana-Champaign

H. Eugene LeMay, Jr.

University of Nevada, Reno

Bruce E. Bursten

University of Tennessee, Knoxville

Catherine J. Murphy

University of Illinois at Urbana-Champaign

Patrick M. Woodward

The Ohio State University

Matthew W. Stoltzfus

The Ohio State University

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Library of Congress Cataloging-In Publication Data

Brown, Theodore L. (Theodore Lawrence), 1928- author.

Chemistry the central science.—Thirteenth edition / Theodore L. Brown, University of Illinois at Urbana-Chanmpaign,

H. Euguene LeMay, Jr., University of Nevada, Reno, Bruce E. Bursten, University of Tennessee, Knoxville,

Catherine J. Murphy, University of Illinois at Urbana-Chanmpaign, Patrick M. Woodward, The Ohio State University,

Matthew W. Stoltzfus, The Ohio State University.

pages cm

Includes index.

ISBN-13: 978-0-321-91041-7

ISBN-10: 0-321-91041-9

1. Chemistry--Textbooks. I. Title.

QD31.3.B765 2014

540—dc23 2013036724

1 2 3 4 5 6 7 8 9 10—CRK— 17 16 15 14

Student Edition: 0-321-91041-9 / 978-0-321-91041-7

www.pearsonhighered.com

Instructor’s Resource Copy: 0-321-96239-7 / 978-0-321-96239-3

To our students,

whose enthusiasm and curiosity have often inspired us,

and whose questions and suggestions have sometimes taught us.

Brief Contents

Preface  xx

1 Introduction: Matter and Measurement  2

2 Atoms, Molecules, and Ions  40

3 Chemical Reactions and Reaction Stoichiometry  80

4 Reactions in Aqueous Solution  122

5 Thermochemistry  164

6 Electronic Structure of Atoms  212

7 Periodic Properties of the Elements  256

8 Basic Concepts of Chemical Bonding  298

9 Molecular Geometry and Bonding Theories  342

10 Gases  398

11 Liquids and Intermolecular Forces  442

12 Solids and Modern Materials  480

13 Properties of Solutions  530

14 Chemical Kinetics  574

15 Chemical Equilibrium  628

16 Acid–Base Equilibria  670

17 Additional Aspects of Aqueous Equilibria  724

18 Chemistry of the Environment  774

19 Chemical Thermodynamics  812

20 Electrochemistry  856

21 Nuclear Chemistry  908

22 Chemistry of the Nonmetals  952

23 Transition Metals and Coordination Chemistry  996

24 The Chemistry of Life: Organic and Biological Chemistry  1040

Appendices  

A Mathematical Operations  1092

B Properties of Water  1099

C Thermodynamic Quantities for Selected Substances at

298.15 K (25 °C)  1100

D Aqueous Equilibrium Constants  1103

E Standard Reduction Potentials at 25 °C  1105

Answers to Selected Exercises  A-1

Answers to Give It Some Thought  A-31

Answers to Go Figure  A-38

Answers to Selected Practice Exercises  A-44

Glossary  G-1

Photo/Art Credits  P-1

Index  I-1

vi

Contents

Preface  xx

1 Introduction: Matter and Measurement  2

1.1 The Study of Chemistry  2

The Atomic and Molecular Perspective of Chemistry  4 Why Study Chemistry?  5

1.2 Classifications of Matter  6

States of Matter  7 Pure Substances  7

Elements  7 Compounds  8 Mixtures  10 1.3 Properties of Matter  11

Physical and Chemical Changes  12

Separation of Mixtures  13

1.4 Units of Measurement  14

SI Units  15 Length and Mass  17

Temperature  17 Derived SI Units  19

Volume  19 Density  19

1.5 Uncertainty in Measurement  22

Precision and Accuracy  22 Significant

Figures  22 Significant Figures in

Calculations  22

1.6 Dimensional Analysis  27

Using Two or More Conversion Factors  28 Conversions Involving Volume  29

Chapter Summary and Key Terms  32

Learning Outcomes  32

Key Equations  32 Exercises  32 Additional Exercises  37

Chemistry Put to Work Chemistry and the

Chemical Industry  6

A Closer Look The Scientific Method  14

Chemistry Put to Work Chemistry in

the News  20

Strategies in Chemistry Estimating Answers  28 Strategies in Chemistry The Importance of Practice  31

Strategies in Chemistry The Features of This Book  32

2 Atoms, Molecules, and Ions  40

2.1 The Atomic Theory of Matter  42 2.2 The Discovery of Atomic Structure  43 Cathode Rays and Electrons  43

Radioactivity  45 The Nuclear Model of the Atom  46

2.3 The Modern View of Atomic Structure  47 Atomic Numbers, Mass Numbers, and

Isotopes  49

2.4 Atomic Weights  50

The Atomic Mass Scale  50 Atomic Weight  51 2.5 The Periodic Table  52

2.6 Molecules and Molecular

Compounds  56

Molecules and Chemical Formulas  56

Molecular and Empirical Formulas  56

Picturing Molecules  57

2.7 Ions and Ionic Compounds  58

Predicting Ionic Charges  59 Ionic

Compounds  60

2.8 Naming Inorganic Compounds  62 Names and Formulas of Ionic Compounds  62 Names and Formulas of Acids  67 Names and Formulas of Binary Molecular Compounds  68

2.9 Some Simple Organic Compounds  69 Alkanes  69 Some Derivatives of Alkanes  70

Chapter Summary and Key Terms  72

Learning Outcomes  72 Key

Equations  73 Exercises  73

Additional Exercises  78

A Closer Look Basic Forces  49

A Closer Look The Mass Spectrometer  52

A Closer Look What Are Coins Made Of?  54 Chemistry and Life Elements Required by Living Organisms  61

Strategies in Chemistry How to Take a Test  71

vii

viii Contents

3 Chemical Reactions and Reaction

Stoichiometry  80

3.1 Chemical Equations  82

Balancing Equations  82 Indicating the States of Reactants and Products  85

3.2 Simple Patterns of Chemical Reactivity  86 Combination and Decomposition

Reactions  86 Combustion Reactions  89

3.3 Formula Weights  89

Formula and Molecular Weights  90

Percentage Composition from Chemical

Formulas  91

3.4 Avogadro’s Number and the Mole  91 Molar Mass  93 Interconverting Masses

and Moles  95 Interconverting Masses and Numbers of Particles  96

3.5 Empirical Formulas from Analyses  98 Molecular Formulas from Empirical

Formulas  100 Combustion Analysis  101

3.6 Quantitative Information from Balanced Equations  103

3.7 Limiting Reactants  106

Theoretical and Percent Yields  109

Chapter Summary and Key Terms  111

Learning Outcomes  111 Key Equations  112 Exercises  112 Additional Exercises  118

Integrative Exercises  120 Design an

Experiment  120

Strategies in Chemistry Problem Solving  92

Chemistry and Life Glucose Monitoring  95

Strategies in Chemistry Design an

Experiment  110

4 Reactions in Aqueous Solution  122

4.1 General Properties of Aqueous

Solutions  124

Electrolytes and Nonelectrolytes  124 How Compounds Dissolve in Water  125 Strong and Weak Electrolytes  126

4.2 Precipitation Reactions  128

Solubility Guidelines for Ionic

Compounds  129 Exchange (Metathesis) Reactions  130 Ionic Equations and Spectator Ions  131

4.3 Acids, Bases, and Neutralization Reactions  132

Acids  132 Bases  133 Strong and Weak Acids and Bases  133 Identifying Strong and Weak Electrolytes  135 Neutralization Reactions and Salts  135 Neutralization Reactions with Gas Formation  138

4.4 Oxidation–Reduction Reactions  138 Oxidation and Reduction  138 Oxidation Numbers  140 Oxidation of Metals by Acids and Salts  142 The Activity Series  143

4.5 Concentrations of Solutions  146

Molarity  146 Expressing the Concentration of an Electrolyte  147 Interconverting Molarity, Moles, and Volume  148 Dilution  149

4.6 Solution Stoichiometry and Chemical Analysis  151

Titrations  152

Chapter Summary and Key Terms  155

Learning Outcomes  156 Key

Equations  156 Exercises  156

Additional Exercises  161 Integrative

Exercises  161 Design an

Experiment  163

Chemistry Put to Work Antacids  139

Strategies in Chemistry Analyzing Chemical Reactions  146

5 Thermochemistry  164

5.1 Energy  166

Kinetic Energy and Potential Energy  166

Units of Energy  168 System and

Surroundings  169 Transferring Energy: Work and Heat  169

5.2 The First Law of Thermodynamics  170 Internal Energy  171 Relating ∆E to Heat and Work  172 Endothermic and Exothermic Processes  173 State Functions  174

5.3 Enthalpy  175

Pressure–Volume Work  175 Enthalpy Change  177

5.4 Enthalpies of Reaction  179

5.5 Calorimetry  181

Heat Capacity and Specific Heat  181

Constant-Pressure Calorimetry  183

Bomb Calorimetry (Constant-Volume

Calorimetry)  185

5.6 Hess’s Law  187

5.7 Enthalpies of Formation  189

Using Enthalpies of Formation to Calculate Enthalpies of Reaction  192

5.8 Foods and Fuels  194

Foods  194 Fuels  197 Other Energy

Sources  198

Chapter Summary and Key Terms  200 Learning Outcomes  201 Key Equations  202 Exercises  202 Additional Exercises  209 Integrative Exercises  210 Design an

Experiment  211

A Closer Look Energy, Enthalpy, and P–V

Work  178

Strategies in Chemistry Using Enthalpy as a Guide  181

Chemistry and Life The Regulation of Body Temperature  186

Chemistry Put to Work The Scientific and Political Challenges of Biofuels  198

6 Electronic Structure of Atoms  212

6.1 The Wave Nature of Light  214

6.2 Quantized Energy and Photons  216 Hot Objects and the Quantization of Energy  216 The Photoelectric Effect and Photons  217

6.3 Line Spectra and the Bohr Model  219 Line Spectra  219 Bohr’s Model  220

The Energy States of the Hydrogen Atom  221 Limitations of the Bohr Model  223

6.4 The Wave Behavior of Matter  223 The Uncertainty Principle  225

6.5 Quantum Mechanics and Atomic Orbitals  226

Contents ix

Orbitals and Quantum Numbers  228

6.6 Representations of Orbitals  230

The s Orbitals  230 The p Orbitals  233

The d and f Orbitals  233

6.7 Many-Electron Atoms  234

Orbitals and Their Energies  234 Electron Spin and the Pauli Exclusion Principle  235

6.8 Electron Configurations  237

Hund’s Rule  237 Condensed Electron

Configurations  239 Transition

Metals  240 The Lanthanides and

Actinides  240

6.9 Electron Configurations and the Periodic Table  241

Anomalous Electron Configurations  245

Chapter Summary and Key Terms  246

Learning Outcomes  247 Key Equations  247 Exercises  248 Additional Exercises  252 Integrative Exercises  255 Design an

Experiment  255

A Closer Look Measurement and the Uncertainty Principle  225

A Closer Look Thought Experiments and

Schrödinger’s Cat  227

A Closer Look Probability Density and Radial Probability Functions  232

Chemistry and Life Nuclear Spin and Magnetic Resonance Imaging  236

7 Periodic Properties of the Elements  256

7.1 Development of the Periodic

Table  258

7.2 Effective Nuclear Charge  259

7.3 Sizes of Atoms and Ions  262

Periodic Trends in Atomic Radii  264 Periodic Trends in Ionic Radii  265

7.4 Ionization Energy  268

Variations in Successive Ionization

Energies  268 Periodic Trends in First

Ionization Energies  268 Electron

Configurations of Ions  271

7.5 Electron Affinity  272

7.6 Metals, Nonmetals, and

Metalloids  273

Metals  274 Nonmetals  276 Metalloids  277

x Contents

7.7 Trends for Group 1A and Group 2A Metals  278

Group 1A: The Alkali Metals  278 Group 2A: The Alkaline Earth Metals  281

7.8 Trends for Selected Nonmetals  282 Hydrogen  282 Group 6A: The Oxygen

Group  283 Group 7A: The Halogens  284

Group 8A: The Noble Gases  286

Chapter Summary and Key Terms  288

Learning Outcomes  289 Key Equations  289 Exercises  289 Additional Exercises  294

Integrative Exercises  296 Design an

Experiment  297

A Closer Look Effective Nuclear Charge  261

Chemistry Put to Work Ionic Size and

Lithium-Ion Batteries  267

Chemistry and Life The Improbable Development of Lithium Drugs  281

8 Basic Concepts of Chemical Bonding  298

8.1 Lewis Symbols and the Octet Rule  300 The Octet Rule  300

8.2 Ionic Bonding  301

Energetics of Ionic Bond Formation  302

Electron Configurations of Ions of the s- and p-Block Elements  305 Transition Metal

Ions  306

8.3 Covalent Bonding  306

Lewis Structures  307 Multiple Bonds  308 8.4 Bond Polarity and Electronegativity  309 Electronegativity  309 Electronegativity and Bond Polarity  310 Dipole Moments  311

Differentiating Ionic and Covalent Bonding  314 8.5 Drawing Lewis Structures  315

Formal Charge and Alternative Lewis

Structures  317

8.6 Resonance Structures  320

Resonance in Benzene  322

8.7 Exceptions to the Octet Rule  322

Odd Number of Electrons  323 Less Than an Octet of Valence Electrons  323 More Than an Octet of Valence Electrons  324

8.8 Strengths and Lengths of Covalent Bonds  325

Bond Enthalpies and the Enthalpies of

Reactions  327 Bond Enthalpy and Bond Length  329

Chapter Summary and Key Terms  332

Learning Outcomes  333 Key Equations  333 Exercises  333 Additional Exercises  338

Integrative Exercises  340 Design an

Experiment  341

A Closer Look Calculation of Lattice Energies: The Born–Haber Cycle  304

A Closer Look Oxidation Numbers, Formal Charges, and Actual Partial Charges  319

Chemistry Put to Work Explosives and Alfred Nobel  330

9 Molecular Geometry and Bonding

Theories  342

9.1 Molecular Shapes  344

9.2 The Vsepr Model  347

Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles  351 Molecules with Expanded Valence Shells  352 Shapes of Larger Molecules  355

9.3 Molecular Shape and Molecular Polarity  356

9.4 Covalent Bonding and Orbital Overlap  358 9.5 Hybrid Orbitals  359

sp Hybrid Orbitals  360 sp² and sp³ Hybrid Orbitals  361 Hypervalent Molecules  362 Hybrid Orbital Summary  364

9.6 Multiple Bonds  365

Resonance Structures, Delocalization, and p Bonding  368 General Conclusions about s and p Bonding  372

9.7 Molecular Orbitals  373

Molecular Orbitals of the Hydrogen

Molecule  373 Bond Order  375

9.8 Period 2 Diatomic Molecules  376

Molecular Orbitals for Li₂ and Be₂  377

Molecular Orbitals from 2p Atomic

Orbitals  377 Electron Configurations for B₂ through Ne₂  381 Electron Configurations and Molecular Properties  383 Heteronuclear Diatomic Molecules  384

Chapter Summary and Key Terms  386

Learning Outcomes  387 Key Equations  388 Exercises  388 Additional Exercises  393

Integrative Exercises  396 Design an

Experiment  397

Chemistry and Life The Chemistry of Vision  372 A Closer Look Phases in Atomic and Molecular Orbitals  379

Chemistry Put to Work Orbitals and Energy  385

10 Gases  398

10.1 Characteristics of Gases  400

10.2 Pressure  401

Atmospheric Pressure and the Barometer  401 10.3 The Gas Laws  404

The Pressure–Volume Relationship: Boyle’s Law  404 The Temperature–Volume

Relationship: Charles’s Law  406 The

Quantity–Volume Relationship: Avogadro’s Law  406

10.4 The Ideal-Gas Equation  408

Relating the Ideal-Gas Equation and the Gas Laws  410

10.5 Further Applications of the Ideal-Gas Equation  412

Gas Densities and Molar Mass  413 Volumes of Gases in Chemical Reactions  414

10.6 Gas Mixtures and Partial

Pressures  415

Partial Pressures and Mole Fractions  417

10.7 The Kinetic-Molecular Theory of Gases  418

Distributions of Molecular Speed  419

Application of Kinetic-Molecular Theory to the Gas Laws  420

10.8 Molecular Effusion and Diffusion  421 Graham’s Law of Effusion  423 Diffusion and Mean Free Path  424

10.9 Real Gases: Deviations from Ideal Behavior  426

The van der Waals Equation  428

Chapter Summary and Key Terms  431

Learning Outcomes  431 Key Equations  432

Contents xi

Exercises  432 Additional Exercises  438

Integrative Exercises  440 Design an

Experiment  441

Strategies in Chemistry Calculations Involving Many Variables  410

A Closer Look The Ideal-Gas Equation  421

Chemistry Put to Work Gas Separations  425

11 Liquids and

Intermolecular

Forces  442

11.1 A Molecular Comparison of Gases, Liquids, and Solids  444

11.2 Intermolecular Forces  446

Dispersion Forces  447 Dipole–Dipole

Forces  448 Hydrogen Bonding  449

Ion–Dipole Forces  452 Comparing

Intermolecular Forces  452

11.3 Select Properties of Liquids  455

Viscosity  455 Surface Tension  456 Capillary Action  456

11.4 Phase Changes  457

Energy Changes Accompanying Phase

Changes  457 Heating Curves  459 Critical Temperature and Pressure  460

11.5 Vapor Pressure  461

Volatility, Vapor Pressure, and

Temperature  462 Vapor Pressure and Boiling Point  463

11.6 Phase Diagrams  464

The Phase Diagrams of H₂O and CO₂  465

11.7 Liquid Crystals  467

Types of Liquid Crystals  467

Chapter Summary and Key Terms  470

Learning Outcomes  471 Exercises  471

Additional Exercises  477 Integrative

Exercises  478 Design an

Experiment  479

Chemistry Put to Work Ionic

Liquids  454

A Closer Look The Clausius–Clapeyron

Equation  463

xii Contents

12 Solids and Modern Materials  480

12.1 Classification of Solids  480

12.2 Structures of Solids  482

Crystalline and Amorphous Solids  482 Unit Cells and Crystal Lattices  483 Filling the Unit Cell  485

12.3 Metallic Solids  486

The Structures of Metallic Solids  487 Close Packing  488 Alloys  491

12.4 Metallic Bonding  494

Electron-Sea Model  494 Molecular–Orbital Model  495

12.5 Ionic Solids  498

Structures of Ionic Solids  498

12.6 Molecular Solids  502

12.7 Covalent-Network Solids  503

Semiconductors  504 Semiconductor

Doping  506

12.8 Polymers  507

Making Polymers  509 Structure and Physical Properties of Polymers  511

12.9 Nanomaterials  514

Semiconductors on the Nanoscale  514 Metals on the Nanoscale  515 Carbons on the

Nanoscale  516

Chapter Summary and Key Terms  519

Learning Outcomes  520 Key Equation  520 Exercises  521 Additional Exercises  527

Integrative Exercises  528 Design an

Experiment  529

A Closer Look X-ray Diffraction  486

Chemistry Put to Work Alloys of Gold  494

Chemistry Put to Work Solid-State

Lighting  508

Chemistry Put to Work Recycling

Plastics  511

13 Properties of

Solutions  530

13.1 The Solution Process  530

The Natural Tendency toward Mixing  532

The Effect of Intermolecular Forces on Solution Formation  532 Energetics of Solution

Formation  533 Solution Formation and

Chemical Reactions  535

13.2 Saturated Solutions and Solubility  536 13.3 Factors Affecting Solubility  538

Solute–Solvent Interactions  538 Pressure Effects  541 Temperature Effects  543

13.4 Expressing Solution Concentration  544 Mass Percentage, ppm, and ppb  544 Mole Fraction, Molarity, and Molality  545

Converting Concentration Units  547

13.5 Colligative Properties  548

Vapor-Pressure Lowering  548 Boiling-Point Elevation  551 Freezing-Point Depression  552 Osmosis  554 Determination of Molar Mass from Colligative Properties  557

13.6 Colloids  559

Hydrophilic and Hydrophobic Colloids  560 Colloidal Motion in Liquids  562

Chapter Summary and Key Terms  564

Learning Outcomes  565 Key Equations  565 Exercises  566 Additional Exercises  571

Integrative Exercises  572 Design an

Experiment  573

Chemistry and Life Fat-Soluble and Water-Soluble Vitamins  539

Chemistry and Life Blood Gases and Deep-Sea Diving  544

A Closer Look Ideal Solutions with Two or More Volatile Components  550

A Closer Look The Van’t Hoff Factor  558

Chemistry and Life Sickle-Cell Anemia  562

Contents xiii

14 Chemical Kinetics  574

15 Chemical

14.1 Factors that Affect Reaction Rates  576 14.2 Reaction Rates  577

Change of Rate with Time  579 Instantaneous Rate  579 Reaction Rates and

Stoichiometry  580

14.3 Concentration and Rate Laws  581 Reaction Orders: The Exponents in the

Rate Law  584 Magnitudes and Units of Rate Constants  585 Using Initial Rates to Determine Rate Laws  586

14.4 The Change of Concentration with Time  587

First-Order Reactions  587 Second-Order Reactions  589 Zero-Order Reactions  591 Half-Life  591

14.5 Temperature and Rate  593

The Collision Model  593 The Orientation Factor  594 Activation Energy  594 The Arrhenius Equation  596 Determining the Activation Energy  597

14.6 Reaction Mechanisms  599

Elementary Reactions  599 Multistep

Mechanisms  600 Rate Laws for Elementary Reactions  601 The Rate-Determining Step for a Multistep Mechanism  602 Mechanisms with a Slow Initial Step  603 Mechanisms with a Fast Initial Step  604

14.7 Catalysis  606

Homogeneous Catalysis  607 Heterogeneous Catalysis  608 Enzymes  609

Chapter Summary and Key Terms  614

Learning Outcomes  614 Key Equations  615 Exercises  615 Additional Exercises  624 Integrative Exercises  626 Design an

Experiment  627

A Closer Look Using Spectroscopic Methods to Measure Reaction Rates: Beer’s Law  582

Chemistry Put to Work Methyl Bromide in the Atmosphere  592

Chemistry Put to Work Catalytic Converters  610 Chemistry and Life Nitrogen Fixation and

Nitrogenase  612

Equilibrium  628

15.1 The Concept of Equilibrium  630

15.2 The Equilibrium Constant  632

Evaluating K_c  634 Equilibrium Constants in Terms of Pressure, K_p  635 Equilibrium Constants and Units  636

15.3 Understanding and Working with Equilibrium Constants  637

The Magnitude of Equilibrium Constants  637 The Direction of the Chemical Equation

and K  639 Relating Chemical Equation

Stoichiometry and Equilibrium Constants  639 15.4 Heterogeneous Equilibria  641

15.5 Calculating Equilibrium Constants  644 15.6 Applications of Equilibrium Constants  646 Predicting the Direction of Reaction  646

Calculating Equilibrium Concentrations  648 15.7 Le Châtelier’s Principle  650

Change in Reactant or Product

Concentration  651 Effects of Volume and Pressure Changes  652 Effect of Temperature Changes  654 The Effect of Catalysts  657

Chapter Summary and Key Terms  660

Learning Outcomes  660 Key Equations  661 Exercises  661 Additional Exercises  666

Integrative Exercises  668 Design an

Experiment  669

Chemistry Put to Work The Haber Process  633 Chemistry Put to Work Controlling Nitric Oxide Emissions  659

16 Acid–Base Equilibria  670

16.1 Acids and Bases: A Brief Review  672 16.2 BrØnsted–Lowry Acids and Bases  673

xiv Contents

The H⁺ Ion in Water  673 Proton-Transfer Reactions  673 Conjugate Acid–Base Pairs  674 Relative Strengths of Acids and Bases  676

16.3 The Autoionization of Water  678 The Ion Product of Water  679

16.4 The pH Scale  680

pOH and Other “p” Scales  682 Measuring pH  683

16.5 Strong Acids and Bases  684

Strong Acids  684 Strong Bases  685

16.6 Weak Acids  686

Calculating K_a from pH  688 Percent

Ionization  689 Using K_a to Calculate pH  690 Polyprotic Acids  694

16.7 Weak Bases  696

Types of Weak Bases  698

16.8 Relationship between K_a and K_b  699 16.9 Acid–Base Properties of Salt Solutions  702 An Anion’s Ability to React with Water  702 A Cation’s Ability to React with Water  702

Combined Effect of Cation and Anion in

Solution  704

16.10 Acid–Base Behavior and Chemical Structure  705

Factors That Affect Acid Strength  705 Binary Acids  706 Oxyacids  707 Carboxylic

Acids  709

16.11 Lewis Acids and Bases  710

Chapter Summary and Key Terms  713

Learning Outcomes  714 Key Equations  714 Exercises  715 Additional Exercises  720

Integrative Exercises  722 Design an

Experiment  723

Chemistry Put to Work Amines and Amine

Hydrochlorides  701

Chemistry and Life The Amphiprotic Behavior of Amino Acids  709

17 Additional Aspects of Aqueous Equilibria  724

17.1 The Common-Ion Effect  726

17.2 Buffers  729

Composition and Action of Buffers  729

Calculating the pH of a Buffer  731 Buffer Capacity and pH Range  734 Addition of

Strong Acids or Bases to Buffers  735

17.3 Acid–Base Titrations  738

Strong Acid–Strong Base Titrations  738 Weak Acid–Strong Base Titrations  740 Titrating with an Acid–Base Indicator  744 Titrations of Polyprotic Acids  746

17.4 Solubility Equilibria  748

The Solubility-Product Constant, K_sp  748

Solubility and K_sp  749

17.5 Factors That Affect Solubility  751 Common-Ion Effect  751 Solubility and

pH  753 Formation of Complex Ions  756

Amphoterism  758

17.6 Precipitation and Separation of Ions  759 Selective Precipitation of Ions  760

17.7 Qualitative Analysis for Metallic Elements  762

Chapter Summary and Key Terms  765

Learning Outcomes  765 Key Equations  766 Exercises  766 Additional Exercises  771

Integrative Exercises  772 Design an

Experiment  773

Chemistry and Life Blood as a Buffered

Solution  737

A Closer Look Limitations of Solubility

Products  751

Chemistry and Life Ocean Acidification  753

Chemistry and Life Tooth Decay and

Fluoridation  755

18 Chemistry of the Environment  774

18.1 Earth’s Atmosphere  776

Composition of the Atmosphere  776

Photochemical Reactions in the

Atmosphere  778 Ozone in the

Stratosphere  780

18.2 Human Activities and Earth’s

Atmosphere  782

The Ozone Layer and Its Depletion  782 Sulfur Compounds and Acid Rain  784 Nitrogen Oxides and Photochemical Smog  786

Greenhouse Gases: Water Vapor, Carbon

Dioxide, and Climate  787

18.3 Earth’s Water  791

The Global Water Cycle  791 Salt Water:

Earth’s Oceans and Seas  792 Freshwater and Groundwater  792

18.4 Human Activities and Water Quality  794 Dissolved Oxygen and Water Quality  794 Water Purification: Desalination  795 Water Purification: Municipal Treatment  796

18.5 Green Chemistry  798

Supercritical Solvents  800 Greener Reagents and Processes  800

Chapter Summary and Key Terms  803

Learning Outcomes  803 Exercises  804

Additional Exercises  808 Integrative

Exercises  809 Design an Experiment  811

A Closer Look Other Greenhouse Gases  790

A Closer Look The Ogallala Aquifer—A Shrinking Resource  794

A Closer Look Fracking and Water Quality  797

19 Chemical

Thermodynamics  812

19.1 Spontaneous Processes  814

Seeking a Criterion for Spontaneity  816

Reversible and Irreversible Processes  816 19.2 Entropy and the Second Law of Thermodynamics  818

The Relationship between Entropy and

Heat  818 ∆S for Phase Changes  819 The Second Law of Thermodynamics  820

19.3 The Molecular Interpretation of Entropy and the Third Law of

Thermodynamics  821

Expansion of a Gas at the Molecular Level  821 Boltzmann’s Equation and Microstates  823 Molecular Motions and Energy  824 Making Qualitative Predictions about ∆S   825 The Third Law of Thermodynamics  827

19.4 Entropy Changes in Chemical

Reactions  828

Entropy Changes in the Surroundings  830 19.5 Gibbs Free Energy  831

Standard Free Energy of Formation  834

19.6 Free Energy and Temperature  836 19.7 Free Energy and the Equilibrium Constant  838

Free Energy under Nonstandard

Conditions  838 Relationship between ∆G° and K  840

Chapter Summary and Key Terms  844

Contents xv

Learning Outcomes  844 Key Equations  845 Exercises  845 Additional Exercises  851

Integrative Exercises  853 Design an

Experiment  855

A Closer Look The Entropy Change When a Gas Expands Isothermally  820

Chemistry and Life Entropy and Human

Society  828

A Closer Look What’s “Free” about Free Energy?  836 Chemistry and Life Driving Nonspontaneous

Reactions: Coupling Reactions  842

20 Electrochemistry  856

20.1 Oxidation States and Oxidation–Reduction Reactions  858

20.2 Balancing Redox Equations  860

Half-Reactions  860 Balancing Equations by the Method of Half-Reactions  860 Balancing Equations for Reactions Occurring in Basic

Solution  863

20.3 Voltaic Cells  865

20.4 Cell Potentials Under Standard

Conditions  868

Standard Reduction Potentials  869 Strengths of Oxidizing and Reducing Agents  874

20.5 Free Energy and Redox Reactions  876 Emf, Free Energy, and the Equilibrium

Constant  877

20.6 Cell Potentials Under Nonstandard Conditions  880

The Nernst Equation  880 Concentration

Cells  882

20.7 Batteries and Fuel Cells  886

Lead–Acid Battery  886 Alkaline Battery  887 Nickel–Cadmium and Nickel–Metal Hydride Batteries  887 Lithium-Ion Batteries  887

Hydrogen Fuel Cells  889

20.8 Corrosion  891

Corrosion of Iron (Rusting)  891 Preventing Corrosion of Iron  892

20.9 Electrolysis  893

Quantitative Aspects of Electrolysis  894

Chapter Summary and Key Terms  897

Learning Outcomes  898 Key Equations  899 Exercises  899 Additional Exercises  905

Integrative Exercises  907 Design an

Experiment  907

xvi Contents

A Closer Look Electrical Work  879

Chemistry and Life Heartbeats and

Electrocardiography  884

Chemistry Put to Work Batteries for Hybrid and Electric Vehicles  889

Chemistry Put to Work Electrometallurgy of

Aluminum  895

21 Nuclear Chemistry  908

21.1 Radioactivity and Nuclear Equations  910 Nuclear Equations  911 Types of Radioactive Decay  912

21.2 Patterns of Nuclear Stability  914

Neutron-to-Proton Ratio  914 Radioactive Decay Chains  916 Further Observations  916 21.3 Nuclear Transmutations  918

Accelerating Charged Particles  918 Reactions Involving Neutrons  919 Transuranium

Elements  920

21.4 Rates of Radioactive Decay  920

Radiometric Dating  921 Calculations Based on Half-Life  923

21.5 Detection of Radioactivity  926

Radiotracers  927

21.6 Energy Changes in Nuclear Reactions  929 Nuclear Binding Energies  930

21.7 Nuclear Power: Fission  932

Nuclear Reactors  934 Nuclear Waste  936 21.8 Nuclear Power: Fusion  937

21.9 Radiation in the Environment and Living Systems  938

Radiation Doses  940 Radon  942

Chapter Summary and Key Terms  944

Learning Outcomes  945 Key Equations  945 Exercises  946 Additional Exercises  949

Integrative Exercises  951 Design an

Experiment  951

Chemistry and Life Medical Applications of

Radiotracers  928

A Closer Look The Dawning of the Nuclear

Age  934

A Closer Look Nuclear Synthesis of the

Elements  939

Chemistry and Life Radiation Therapy  943

22 Chemistry of the Nonmetals  952

22.1 Periodic Trends and Chemical

Reactions  952

Chemical Reactions  955

22.2 Hydrogen  956

Isotopes of Hydrogen  956 Properties of

Hydrogen  957 Production of Hydrogen  958 Uses of Hydrogen  959 Binary Hydrogen

Compounds  959

22.3 Group 8A: The Noble Gases  960

Noble-Gas Compounds  961

22.4 Group 7A: The Halogens  962

Properties and Production of the Halogens  962 Uses of the Halogens  964 The Hydrogen Halides  964 Interhalogen Compounds  965 Oxyacids and Oxyanions  966

22.5 Oxygen  966

Properties of Oxygen  967 Production of

Oxygen  967 Uses of Oxygen  967

Ozone  967 Oxides  968 Peroxides and

Superoxides  969

22.6 The Other Group 6A Elements: S, Se, Te, and Po  970

General Characteristics of the Group 6A

Elements  970 Occurrence and Production of S, Se, and Te  970 Properties and Uses of Sulfur, Selenium, and Tellurium  971

Sulfides  971 Oxides, Oxyacids, and

Oxyanions of Sulfur  971

22.7 Nitrogen  973

Properties of Nitrogen  973 Production and Uses of Nitrogen  973 Hydrogen Compounds of Nitrogen  973 Oxides and Oxyacids of Nitrogen  975

22.8 The Other Group 5A Elements: P, As, Sb, and Bi  977

General Characteristics of the Group 5A

Elements  977 Occurrence, Isolation, and Properties of Phosphorus  977 Phosphorus Halides  978 Oxy Compounds of

Phosphorus  978

22.9 Carbon  980

Elemental Forms of Carbon  980 Oxides

of Carbon  981 Carbonic Acid and

Carbonates  983 Carbides  983

22.10 The Other Group 4A Elements: Si, Ge, Sn, and Pb  984

General Characteristics of the Group 4A

Elements  984 Occurrence and Preparation of Silicon  984 Silicates  985 Glass  986

Silicones  987

22.11 Boron  987

Chapter Summary and Key Terms  989

Learning Outcomes  990 Exercises  990

Additional Exercises  994 Integrative

Exercises  994 Design an Experiment  995

A Closer Look The Hydrogen Economy  958

Chemistry and Life Nitroglycerin, Nitric Oxide, and Heart Disease  976

Chemistry and Life Arsenic in

Drinking Water  980

Chemistry Put to Work Carbon Fibers and

Composites  982

23 Transition Metals and Coordination

Chemistry  996

23.1 The Transition Metals  998

Physical Properties  998

Electron Configurations and Oxidation

States  999 Magnetism  1001

23.2 Transition-Metal Complexes  1002 The Development of Coordination Chemistry: Werner’s Theory  1003 The Metal–Ligand Bond  1005 Charges, Coordination Numbers, and Geometries  1006

23.3 Common Ligands in Coordination Chemistry  1007

Metals and Chelates in Living Systems  1009 23.4 Nomenclature and Isomerism in Coordination Chemistry  1012

Isomerism  1014 Structural Isomerism  1014 Stereoisomerism  1015

23.5 Color and Magnetism in Coordination Chemistry  1019

Color  1019 Magnetism of Coordination

Compounds  1021

23.6 Crystal-Field Theory  1021

Contents xvii

Electron Configurations in Octahedral

Complexes  1024 Tetrahedral and Square Planar Complexes  1026

Chapter Summary and Key Terms  1030

Learning Outcomes  1031 Exercises  1031

Additional Exercises  1035 Integrative

Exercises  1037 Design an Experiment  1039 A Closer Look Entropy and the Chelate

Effect  1010

Chemistry and Life The Battle for Iron in Living Systems  1011

A Closer Look Charge-Transfer Color  1028

24 The Chemistry of Life: Organic and Biological Chemistry  1040

24.1 General Characteristics of Organic Molecules  1042

The Structures of Organic Molecules  1042

The Stabilities of Organic Substances  1043 Solubility and Acid–Base Properties of Organic Substances  1042

24.2 Introduction to Hydrocarbons  1044 Structures of Alkanes  1045 Structural

Isomers  1045 Nomenclature of Alkanes  1046 Cycloalkanes  1049 Reactions of

Alkanes  1049

24.3 Alkenes, Alkynes, and Aromatic Hydrocarbons  1050

Alkenes  1051 Alkynes  1053 Addition

Reactions of Alkenes and Alkynes  1054

Aromatic Hydrocarbons  1056 Stabilization of p Electrons by Delocalization  1056

Substitution Reactions  1057

24.4 Organic Functional Groups  1058

Alcohols  1058 Ethers  1061 Aldehydes

and Ketones  1061 Carboxylic Acids and

Esters  1062 Amines and Amides  1066

24.5 Chirality in Organic

Chemistry  1067

24.6 Introduction to Biochemistry  1067 24.7 Proteins  1068

Amino Acids  1068 Polypeptides and

Proteins  1070 Protein Structure  1071

xviii Contents

24.8 Carbohydrates  1073

Disaccharides  1074 Polysaccharides  1075 24.9 Lipids  1076

Fats  1076 Phospholipids  1077

24.10 Nucleic Acids  1077

Chapter Summary and Key Terms  1082 Learning Outcomes  1083 Exercises  1083 Additional Exercises  1089

Integrative Exercises  1090

Design an Experiment  1091

Chemistry Put to Work Gasoline  1050

A Closer Look Mechanism of Addition

Reactions  1055

Strategies in Chemistry What Now?  1081

Appendices  

A Mathematical Operations  1092

B Properties of Water  1099

C Thermodynamic Quantities

for Selected Substances AT 298.15 K (25 °C)  1100

D Aqueous Equilibrium Constants  1103 E Standard Reduction Potentials at 25 °C  1105

Answers to Selected Exercises  A-1

Answers to Give It Some Thought  A-31 Answers to Go Figure  A-38

Answers to Selected Practice Exercises  A-44 Glossary  G-1

Photo/Art Credits  P-1

Index  I-1

Chemical Applications and Essays

Chemistry Put to Work  

Chemistry and the Chemical Industry  6 Chemistry in the News   20

Antacids  139

The Scientific and Political Challenges of Biofuels   198 Ionic Size and Lithium-Ion Batteries  267

Explosives and Alfred Nobel   330

Orbitals and Energy   385

Gas Separations   425

Ionic Liquids   454

Alloys of Gold   494

Solid-State Lighting   508

Recycling Plastics   511

Methyl Bromide in the Atmosphere  592

Catalytic Converters   610

The Haber Process   633

Controlling Nitric Oxide Emissions  659

Amines and Amine Hydrochlorides   701

Batteries for Hybrid and Electric Vehicles   889 Electrometallurgy of Aluminum   895

Carbon Fibers and Composites   982

Gasoline  1050

A Closer Look  

The Scientific Method   14

Basic Forces   49

The Mass Spectrometer   52

What Are Coins Made Of?  54

Energy, Enthalpy, and P–V Work   178

Measurement and the Uncertainty Principle   225 Thought Experiments and Schrödinger’s Cat  226 Probability Density and Radial Probability Functions   232 Effective Nuclear Charge   261

Calculation of Lattice Energies: The Born–Haber Cycle  304 Oxidation Numbers, Formal Charges, and Actual Partial Charges   319

Phases in Atomic and Molecular Orbitals   379 The Ideal-Gas Equation   421

The Clausius–Clapeyron Equation   463

X-ray Diffraction   486

Ideal Solutions with Two or More Volatile Components   550 The Van’t Hoff Factor  558

Using Spectroscopic Methods to Measure Reaction Rates: Beer’s Law  582

Limitations of Solubility Products  751

Other Greenhouse Gases   790

The Ogallala Aquifer—A Shrinking Resource  794 Fracking and Water Quality  797

The Entropy Change When a Gas Expands Isothermally   820 What’s “Free” about Free Energy?  836

Electrical Work  879

The Dawning of the Nuclear Age   934

Nuclear Synthesis of the Elements   939

The Hydrogen Economy   958

Entropy and the Chelate Effect   1010

Charge-Transfer Color   1028

Mechanism of Addition Reactions   1055

Chemistry and Life  

Elements Required by Living Organisms   61 Glucose Monitoring  95

The Regulation of Body Temperature   186 Nuclear Spin and Magnetic Resonance Imaging   236 The Improbable Development of Lithium Drugs   281 The Chemistry of Vision   372

Fat-Soluble and Water-Soluble Vitamins   539

Blood Gases and Deep-Sea Diving   544

Sickle-Cell Anemia   562

Nitrogen Fixation and Nitrogenase   612

The Amphiprotic Behavior of Amino Acids   709 Blood as a Buffered Solution   737

Ocean Acidification   753

Tooth Decay and Fluoridation  755

Entropy and Human Society   828

Driving Nonspontaneous Reactions: Coupling Reactions  842 Heartbeats and Electrocardiography  884

Medical Applications of Radiotracers   928

Radiation Therapy   943

Nitroglycerin, Nitric Oxide, and Heart Disease   976 Arsenic in Drinking Water   980

The Battle for Iron in Living Systems   1011

Strategies in Chemistry  

Estimating Answers   28

The Importance of Practice   31 The Features of This Book   32

How to Take a Test  71

Problem Solving   92

Design an Experiment  110

Analyzing Chemical Reactions   146

Using Enthalpy as a Guide   181

Calculations Involving Many Variables   410

What Now?  1081

xix

Preface

To the Instructor

Philosophy

We authors of Chemistry: The Central Science are delighted and honored that you have chosen us as your instructional partners for your general chemistry class. We have all been active researchers who appreciate both the learning and the discovery aspects of the chemical sciences. We have also all taught general chemistry many times. Our varied, wide-ranging experiences have formed the basis of the close collaborations we have enjoyed as coauthors. In writing our book, our focus is on the students: we try to ensure that the text is not only accurate and up-to-date but also clear and readable. We strive to convey the breadth of chemistry and the excitement that scientists experience in making new discoveries that contribute to our understanding of the physical world. We want the student to appreciate that chemistry is not a body of specialized knowledge that is separate from most aspects of modern life, but central to any attempt to address a host of societal concerns, including renewable energy, environmental sustainability, and improved human health.

Publishing the thirteenth edition of this text bespeaks an exceptionally long record of successful textbook writing. We are appreciative of the loyalty and support the book has received over the years, and mindful of our obligation to justify each new edition. We begin our approach to each new edition with an in

tensive author retreat, in which we ask ourselves the deep ques tions that we must answer before we can move forward. What justifies yet another edition? What is changing in the world not only of chemistry, but with respect to science education and the qualities of the students we serve? The answer lies only partly in the changing face of chemistry itself. The introduction of many new technologies has changed the landscape in the teach ing of sciences at all levels. The use of the Internet in accessing information and presenting learning materials has markedly changed the role of the textbook as one element among many tools for student learning. Our challenge as authors is to main tain the text as the primary source of chemical knowledge and practice, while at the same time integrating it with the new ave nues for learning made possible by technology and the Internet. This edition incorporates links to a number of those new meth odologies, including use of the Internet, computer-based class room tools, such as Learning Catalytics™, a cloud-based active learning analytics and assessment system, and web-based tools, ^{particularly MasteringChemistry}®^{, which is continually evolv} ing to provide more effective means of testing and evaluating student performance, while giving the student immediate and ^{helpful feedback. In past versions, MasteringChemistry}_® ^pro vided feedback only on a question level. Now with Knewton enhanced adaptive follow-up assignments, and Dynamic Study ^{Modules, MasteringChemistry}_® ^{continually adapts to each stu} dent, offering a personalized learning experience.

As authors, we want this text to be a central, indispensa ble learning tool for students. Whether as a physical book or in electronic form, it can be carried everywhere and used at any time. It is the one place students can go to obtain the informa tion outside of the classroom needed for learning, skill develop ment, reference, and test preparation. The text, more effectively than any other instrument, provides the depth of coverage and coherent background in modern chemistry that students need to serve their professional interests and, as appropriate, to pre pare for more advanced chemistry courses.

If the text is to be effective in supporting your role as in structor, it must be addressed to the students. We have done our best to keep our writing clear and interesting and the book attractive and well illustrated. The book has numerous in-text study aids for students, including carefully placed descrip tions of problem-solving strategies. We hope that our cumula tive experiences as teachers is evident in our pacing, choice of examples, and the kinds of study aids and motivational tools we have employed. We believe students are more enthusiastic about learning chemistry when they see its importance relative to their own goals and interests; therefore, we have highlighted many important applications of chemistry in everyday life. We hope you make use of this material.

It is our philosophy, as authors, that the text and all the sup plementary materials provided to support its use must work in concert with you, the instructor. A textbook is only as useful to students as the instructor permits it to be. This book is replete with features that can help students learn and that can guide them as they acquire both conceptual understanding and prob lem-solving skills. There is a great deal here for the students to use, too much for all of it to be absorbed by any one student. You will be the guide to the best use of the book. Only with your active help will the students be able to utilize most effectively all that the text and its supplements offer. Students care about grades, of course, and with encouragement they will also be come interested in the subject matter and care about learning. Please consider emphasizing features of the book that can en hance student appreciation of chemistry, such as the Chemistry Put to Work and Chemistry and Life boxes that show how chem istry impacts modern life and its relationship to health and life processes. Learn to use, and urge students to use, the rich online resources available. Emphasize conceptual understanding and place less emphasis on simple manipulative, algorithmic prob lem solving.

What Is New in This Edition?

A great many changes have been made in producing this thir teenth edition. We have continued to improve upon the art program, and new features connected with the art have been introduced. Many figures in the book have undergone modifi cation, and dozens of new figures have been introduced.

xx

A systematic effort has been made to place explanatory la bels directly into figures to guide the student. New designs have been employed to more closely integrate photographic materi als into figures that convey chemical principles.

We have continued to explore means for more clearly and directly addressing the issue of concept learning. It is well es tablished that conceptual misunderstandings, which impede student learning in many areas, are difficult to correct. We have looked for ways to identify and correct misconceptions via the worked examples in the book, and in the accompanying prac tice exercises. Among the more important changes made in the new edition, with this in mind, are:

• A major new feature of this edition is the addition of a second Practice Exercise to accompany each Sample Ex ercise within the chapters. The majority of new Practice Exercises are of the multiple-choice variety, which enable ^{feedback via MasteringChemistry}®^{. The correct answers} to select Practice Exercises are given in an appendix, and guidance for correcting wrong answers is provided in Mas ^{teringChemistry}®^{. The new Practice Exercise feature adds} to the aids provided to students for mastering the concepts advanced in the text and rectifying conceptual misunder standings. The enlarged practice exercise materials also further cement the relationship of the text to the online learning materials. At the same time, they offer a new sup portive learning experience for all students, regardless of ^{whether the MasteringChemistry}_® program is used.

• A second major innovation in this edition is the Design An Experiment feature, which appears as a final exercise in all chapters beginning with Chapter 3, as well as in ^{MasteringChemistry}®^{. The Design an Experiment exercise is} a departure from the usual kinds of end-of-chapter exer

cises in that it is inquiry based, open ended, and tries to stimulate the student to “think like a scientist.” Each exer cise presents the student with a scenario in which vari ous unknowns require investigation. The student is called upon to ponder how experiments might be set up to pro vide answers to particular questions about a system, and/ or test plausible hypotheses that might account for a set of observations. The aim of the Design an Experiment exer cises is to foster critical thinking. We hope that they will be effective in active learning environments, which include classroom-based work and discussions, but they are also suitable for individual student work. There is no one right way to solve these exercises, but we authors offer some ideas in an online Instructor’s Resource Manual, which will include results from class testing and analysis of stu dent responses.

• The Go Figure exercises introduced in the twelfth edition proved to be a popular innovation, and we have expanded on its use. This feature poses a question that students can answer by examining the figure. These questions encour

age students to actually study the figure and understand its primary message. Answers to the Go Figure questions are provided in the back of the text.

• The popular Give It Some Thought (GIST) questions em bedded in the text have been expanded by improvements

Preface xxi

in some of the existing questions and addition of new ones. The answers to all the GIST items are provided in the back of the text.

• New end-of-chapter exercises have been added, and many of those carried over from the twelfth edition have been significantly revised. Analysis of student responses to the ^{twelfth edition questions in MasteringChemistry}_® helped us identify and revise or create new questions, prompt

ing improvements and eliminations of some questions. ^{Additionally, analysis of usage of MasteringChemistry}® _{has enhanced our understanding of the ways in which in} structors and students have used the end-of-chapter and ^{MasteringChemistry}_® materials. This, in turn, has led to additional improvements to the content within the text ^{and in the MasteringChemistry}_® item library. At the end of each chapter, we list the Learning Outcomes that students should be able to perform after studying each section. End-of-chapter exercises, both in the text and in Master ^ingChemistry_® offer ample opportunities for students to assess mastery of learning outcomes. We trust the Learning Outcomes will help you organize your lectures and tests as the course proceeds.

Organization and Contents

The first five chapters give a largely macroscopic, phenomeno logical view of chemistry. The basic concepts introduced—such as nomenclature, stoichiometry, and thermochemistry—provide necessary background for many of the laboratory experiments usually performed in general chemistry. We believe that an early introduction to thermochemistry is desirable because so much of our understanding of chemical processes is based on consid erations of energy changes. Thermochemistry is also important when we come to a discussion of bond enthalpies. We believe we have produced an effective, balanced approach to teaching ther modynamics in general chemistry, as well as providing students with an introduction to some of the global issues involving en ergy production and consumption. It is no easy matter to walk the narrow pathway between—on the one hand—trying to teach too much at too high a level and—on the other hand—resorting to oversimplifications. As with the book as a whole, the emphasis has been on imparting conceptual understanding, as opposed to presenting equations into which students are supposed to plug numbers.

The next four chapters (Chapters 6–9) deal with elec tronic structure and bonding. We have largely retained our presentation of atomic orbitals. For more advanced students, Closer Look boxes in Chapters 6 and 9 highlight radial prob ability functions and the phases of orbitals. Our approach of placing this latter discussion in a Closer Look box in Chapter 9 enables those who wish to cover this topic to do so, while others may wish to bypass it. In treating this topic and others in Chapters 7 and 9, we have materially enhanced the accom panying figures to more effectively bring home their central messages.

In Chapters 10–13, the focus of the text changes to the next level of the organization of matter: examining the states of

xxii Preface

matter. Chapters 10 and 11 deal with gases, liquids, and inter molecular forces, as in earlier editions. Chapter 12 is devoted to solids, presenting an enlarged and more contemporary view of the solid state as well as of modern materials. The chapter provides an opportunity to show how abstract chemical bond ing concepts impact real-world applications. The modular organization of the chapter allows you to tailor your coverage to focus on materials (semiconductors, polymers, nanomaterials, and so forth) that are most relevant to your students and your own interests. Chapter 13 treats the formation and properties

of solutions in much the same manner as the previous edition. The next several chapters examine the factors that determine the speed and extent of chemical reactions: kinetics (Chapter 14), equilibria (Chapters 15–17), thermodynamics (Chapter 19), and electrochemistry (Chapter 20). Also in this section is a chapter on environmental chemistry (Chapter 18), in which the concepts developed in preceding chapters are applied to a discussion of the atmosphere and hydrosphere. This chapter has increasingly come to be focused on green chemistry and the impacts of human activi ties on Earth’s water and atmosphere.

After a discussion of nuclear chemistry (Chapter 21), the book ends with three survey chapters. Chapter 22 deals with nonmetals, Chapter 23 with the chemistry of transition metals, including coordination compounds, and Chapter 24 with the chemistry of organic compounds and elementary biochemical themes. These final four chapters are developed in a parallel fashion and can be covered in any order.

Our chapter sequence provides a fairly standard organ ization, but we recognize that not everyone teaches all the topics in the order we have chosen. We have therefore made sure that instructors can make common changes in teaching sequence with no loss in student comprehension. In particu lar, many instructors prefer to introduce gases (Chapter 10) after stoichiometry (Chapter 3) rather than with states of matter. The chapter on gases has been written to permit this change with no disruption in the flow of material. It is also possible to treat balancing redox equations (Sections 20.1 and 20.2) earlier, after the introduction of redox reactions in Section 4.4. Finally, some instructors like to cover organic chemistry (Chapter 24) right after bonding (Chapters 8 and 9). This, too, is a largely seamless move.

We have brought students into greater contact with de scriptive organic and inorganic chemistry by integrating exam ples throughout the text. You will find pertinent and relevant examples of “real” chemistry woven into all the chapters to il lustrate principles and applications. Some chapters, of course, more directly address the “descriptive” properties of elements and their compounds, especially Chapters 4, 7, 11, 18, and 22–24. We also incorporate descriptive organic and inorganic chemistry in the end-of-chapter exercises.

Changes in This Edition

The What is New in This Edition section on pp. xx–xxi details changes made throughout the new edition. Beyond a mere list ing, however, it is worth dwelling on the general goals we set forth in formulating this new edition. Chemistry: The Central

Science has traditionally been valued for its clarity of writing, its scientific accuracy and currency, its strong end-of-chapter exercises, and its consistency in level of coverage. In making changes, we have made sure not to compromise these charac

teristics, and we have also continued to employ an open, clean design in the layout of the book.

The art program for this thirteenth edition has continued the trajectory set in the twelfth edition: to make greater and more effective use of the figures as learning tools, by drawing the reader more directly into the figure. The art itself has con

tinued to evolve, with modifications of many figures and addi tions or replacements that teach more effectively. The Go Figure feature has been expanded greatly to include a larger number of figures. In the same vein, we have added to the Give it Some Thought feature, which stimulates more thoughtful reading of the text and fosters critical thinking.

We provide a valuable overview of each chapter under the What’s Ahead banner. Concept links ( ) continue to provide easy-to-see cross-references to pertinent material covered ear lier in the text. The essays titled Strategies in Chemistry, which provide advice to students on problem solving and “thinking like a chemist,” continue to be an important feature. For exam ple, the new Strategies in Chemistry essay at the end of Chapter 3 introduces the new Design an Experiment feature and provides a worked out example as guidance.

We have continued to emphasize conceptual exercises in the end-of-chapter exercise materials. The well-received Visu alizing Concepts exercise category has been continued in this edition. These exercises are designed to facilitate concept un derstanding through use of models, graphs, and other visual materials. They precede the regular end-of-chapter exercises and are identified in each case with the relevant chapter section number. A generous selection of Integrative Exercises, which give students the opportunity to solve problems that integrate concepts from the present chapter with those of previous chap ters, is included at the end of each chapter. The importance of integrative problem solving is highlighted by the Sample Integrative Exercise, which ends each chapter beginning with Chapter 4. In general, we have included more conceptual end of-chapter exercises and have made sure that there is a good representation of somewhat more difficult exercises to provide a better mix in terms of topic and level of difficulty. Many of the exercises have been restructured to facilitate their use in Mas ^{teringChemistry}®^{. We have made extensive use of the metadata from student use of MasteringChemistry}_® ^{to analyze end-of} chapter exercises and make appropriate changes, as well as to develop Learning Outcomes for each chapter.

New essays in our well-received Chemistry Put to Work and Chemistry and Life series emphasize world events, scientific discoveries, and medical breakthroughs that bear on topics de veloped in each chapter. We maintain our focus on the positive aspects of chemistry without neglecting the problems that can arise in an increasingly technological world. Our goal is to help students appreciate the real-world perspective of chemistry and the ways in which chemistry affects their lives.

It is perhaps a natural tendency for chemistry text books to grow in length with succeeding editions, but it is

one that we have resisted. There are, nonetheless, many new items in this edition, mostly ones that replace other material considered less pertinent. Here is a list of several significant changes in content:

In Chapter 1, the Closer Look box on the scientific method has been rewritten. The Chemistry Put to Work box, dealing with Chemistry in the News, has been completely rewritten, with items that describe diverse ways in which chemistry intersects with the affairs of modern society. The Chapter Summary and Learning Outcomes sections at the end of the chapter have been rewritten for ease of use by both instructor and student, in this and all chapters in the text. Similarly, the exercises have been thoroughly vetted, modified where this was called for and re placed or added to, here and in all succeeding chapters.

In Chapter 3, graphic elements highlighting the correct ap proach to problem solving have been added to Sample Exercises on calculating an empirical formula from mass percent of the elements present, combustion analysis, and calculating a theo retical yield.

Chapter 5 now presents a more explicit discussion of com bined units of measurement, an improved introduction to en thalpy, and more consistent use of color in art.

Changes in Chapter 6 include a significant revision of the discussion of the energy levels of the hydrogen atom, including greater clarity on absorption versus emission processes. There is also a new Closer Look box on Thought Experiments and Schrödinger’s Cat, which gives students a brief glimpse of some of the philosophical issues in quantum mechanics and also con

nects to the 2012 Nobel Prize in Physics.

In Chapter 7, the emphasis on conceptual thinking was en hanced in several ways: the section on effective nuclear charge was significantly revised to include a classroom-tested analogy, the number of Go Figure features was increased substantially, and new end-of-chapter exercises emphasize critical thinking and understanding concepts. In addition, the Chemistry Put to Work box on lithium-ion batteries was updated and revised to include discussion of current issues in using these batteries. Fi nally, the values of ionic radii were revised to be consistent with a recent research study of the best values for these radii.

In Chapter 9, which is one of the most challenging for students, we continue to refine our presentation based on our classroom experience. Twelve new Go Figure exercises will stim ulate more student thought in a chapter with a large amount of graphic material. The discussion of molecular geometry was made more conceptually oriented. The section on delocalized bonding was completely revised to provide what we believe will be a better introduction that students will find useful in organic chemistry. The Closer Look box on phases in orbitals was re

vamped with improved artwork. We also increased the number of end-of-chapter exercises, especially in the area of molecular orbital theory. The Design an Experiment feature in this chapter gives the students the opportunity to explore color and conju

gated π systems.

Chapter 10 contains a new Sample Exercise that walks the student through the calculations that are needed to understand Torricelli’s barometer. Chapter 11 includes an improved defini tion of hydrogen bonding and updated data for the strengths

Preface xxiii

of intermolecular attractions. Chapter 12 includes the latest up dates to materials chemistry, including plastic electronics. New material on the diffusion and mean free path of colloids in solu tion is added to Chapter 13, making a connection to the diffu sion of gas molecules from Chapter 10.

In Chapter 14, ten new Go Figure exercises have been added to reinforce many of the concepts presented as figures and graphs in the chapter. The Design an Experiment exercise in the chapter connects strongly to the Closer Look box on Beer’s Law, which is often the basis for spectrometric kinetics experi

ments performed in the general chemistry laboratory. The presentation in Chapter 16 was made more closely tied to that in Chapter 15, especially through the use of more initial/ change/equilibrium (ICE) charts. The number of conceptual end-of-chapter exercises, including Visualizing Concepts fea tures, was increased significantly.

Chapter 17 offers improved clarity on how to make buff ers, and when the Henderson–Hasselbalch equation may not be accurate. Chapter 18 has been extensively updated to reflect changes in this rapidly evolving area of chemistry. Two Closer Look boxes have been added; one dealing with the shrinking level of water in the Ogallala aquifer and a second with the po tential environmental consequences of hydraulic fracking. In Chapter 20, the description of Li-ion batteries has been signifi cantly expanded to reflect the growing importance of these bat teries, and a new Chemistry Put to Work box on batteries for hybrid and electric vehicles has been added.

Chapter 21 was updated to reflect some of the current is sues in nuclear chemistry and more commonly used nomencla ture for forms of radiation are now used. Chapter 22 includes an improved discussion of silicates.

In Chapter 23, the section on crystal-field theory (Section 23.6) has undergone considerable revision. The description of how the d-orbital energies of a metal ion split in a tetrahedral crystal field has been expanded to put it on par with our treat

ment of the octahedral geometry, and a new Sample Exercise that effectively integrates the links between color, magnetism, and the spectrochemical series has been added. Chapter 24’s coverage of organic chemistry and biochemistry now includes oxidation–reduction reactions that organic chemists find most relevant.

To the Student

Chemistry: The Central Science, Thirteenth Edition, has been writ ten to introduce you to modern chemistry. As authors, we have, in effect, been engaged by your instructor to help you learn chemistry. Based on the comments of students and instructors who have used this book in its previous editions, we believe that we have done that job well. Of course, we expect the text to continue to evolve through future editions. We invite you to write to tell us what you like about the book so that we will know where we have helped you most. Also, we would like to learn of any shortcomings so that we might further improve the book in subsequent editions. Our ad

dresses are given at the end of the Preface.

xxiv Preface

Advice for Learning and

Studying Chemistry

Learning chemistry requires both the assimilation of many con cepts and the development of analytical skills. In this text, we have provided you with numerous tools to help you succeed in both tasks. If you are going to succeed in your chemistry course, you will have to develop good study habits. Science courses, and chemistry in particular, make different demands on your learn ing skills than do other types of courses. We offer the following tips for success in your study of chemistry:

Don’t fall behind! As the course moves along, new top ics will build on material already presented. If you don’t keep up in your reading and problem solving, you will find it much harder to follow the lectures and discussions on current topics. Experienced teachers know that students who read the relevant sections of the text before coming to a class learn more from the class and retain greater recall. “Cramming” just before an exam has been shown to be an ineffective way to study any subject, chemistry included. So now you know. How important to you, in this competitive world, is a good grade in chemistry?

Focus your study. The amount of information you will be expected to learn can sometimes seem overwhelming. It is essential to recognize those concepts and skills that are par ticularly important. Pay attention to what your instructor is emphasizing. As you work through the Sample Exercises and homework assignments, try to see what general principles and skills they employ. Use the What’s Ahead feature at the begin ning of each chapter to help orient yourself to what is important in each chapter. A single reading of a chapter will simply not be enough for successful learning of chapter concepts and prob lem-solving skills. You will need to go over assigned materials more than once. Don’t skip the Give It Some Thought and Go Figure features, Sample Exercises, and Practice Exercises. They are your guides to whether you are learning the material. They are also good preparation for test-taking. The Learning Out comes and Key Equations at the end of the chapter should help you focus your study.

Keep good lecture notes. Your lecture notes will provide you with a clear and concise record of what your instructor regards as the most important material to learn. Using your lecture notes in conjunction with this text is the best way to de

termine which material to study.

Skim topics in the text before they are covered in lecture. Reviewing a topic before lecture will make it easier for you to take good notes. First read the What’s Ahead points and the end-of-chapter Summary; then quickly read through the chap ter, skipping Sample Exercises and supplemental sections. Pay ing attention to the titles of sections and subsections gives you

a feeling for the scope of topics. Try to avoid thinking that you must learn and understand everything right away. You need to do a certain amount of preparation before lecture. More than ever, instructors are using the lecture pe riod not simply as a one-way channel of communication from teacher to student. Rather, they expect students to come to class ready to work on problem solving and critical thinking. Com ing to class unprepared is not a good idea for any lecture envi ronment, but it certainly is not an option for an active learning classroom if you aim to do well in the course.

After lecture, carefully read the topics covered in class. As you read, pay attention to the concepts presented and to the application of these concepts in the Sample Exercises. Once you think you understand a Sample Exercise, test your understand ing by working the accompanying Practice Exercise.

Learn the language of chemistry. As you study chemis try, you will encounter many new words. It is important to pay attention to these words and to know their meanings or the entities to which they refer. Knowing how to identify chemi cal substances from their names is an important skill; it can help you avoid painful mistakes on examinations. For example, “chlorine” and “chloride” refer to very different things.

Attempt the assigned end-of-chapter exercises. Work ing the exercises selected by your instructor provides necessary practice in recalling and using the essential ideas of the chapter. You cannot learn merely by observing; you must be a partici pant. In particular, try to resist checking the Student Solutions Manual (if you have one) until you have made a sincere effort to solve the exercise yourself. If you get stuck on an exercise, however, get help from your instructor, your teaching assistant, or another student. Spending more than 20 minutes on a single exercise is rarely effective unless you know that it is particularly challenging.

Learn to think like a scientist. This book is written by sci entists who love chemistry. We encourage you to develop your critical thinking skills by taking advantage of new features in this edition, such as exercises that focus on conceptual learning, and the Design an Experiment exercises.

Use online resources. Some things are more easily learned by discovery, and others are best shown in three dimensions. ^{If your instructor has included MasteringChemistry}_® with your book, take advantage of the unique tools it provides to get the most out of your time in chemistry.

The bottom line is to work hard, study effectively, and use the tools available to you, including this textbook. We want to help you learn more about the world of chemistry and why chemistry is the central science. If you really learn chemistry, you can be the life of the party, impress your friends and par

ents, and … well, also pass the course with a good grade.

Acknowledgments

The production of a textbook is a team effort requiring the in volvement of many people besides the authors who contributed hard work and talent to bring this edition to life. Although their names don’t appear on the cover of the book, their creativity, time, and support have been instrumental in all stages of its de velopment and production.

Each of us has benefited greatly from discussions with colleagues and from correspondence with instructors and stu

Thirteenth Edition Reviewers

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Preface xxv

dents both here and abroad. Colleagues have also helped im mensely by reviewing our materials, sharing their insights, and providing suggestions for improvements. On this edition, we were particularly blessed with an exceptional group of accuracy checkers who read through our materials looking for both tech nical inaccuracies and typographical errors.

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xxvi Preface

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Preface xxvii

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xxviii Preface

We would also like to express our gratitude to our many team members at Pearson whose hard work, imagination, and com mitment have contributed so greatly to the final form of this edition: Terry Haugen, our senior editor, who has brought en ergy and imagination to this edition as he has to earlier ones; Chris Hess, our chemistry editor, for many fresh ideas and his unflagging enthusiasm, continuous encouragement, and sup port; Jennifer Hart, Director of Development, who has brought her experience and insight to oversight of the entire project; Jessica Moro, our project editor, who very effectively coordinat ed the scheduling and tracked the multidimensional deadlines that come with a project of this magnitude; Jonathan Cottrell our marketing manager, for his energy, enthusiasm, and crea tive promotion of our text; Carol Pritchard-Martinez, our development editor, whose depth of experience, good judgment, and careful attention to detail were invaluable to this revision,

especially in keeping us on task in terms of consistency and stu dent understanding; Donna, our copy editor, for her keen eye; Beth Sweeten, our project manager, and Gina Cheselka, who managed the complex responsibilities of bringing the design, photos, artwork, and writing together with efficiency and good

cheer. The Pearson team is a first-class operation. There are many others who also deserve special recogni tion, including the following: Greg Johnson, our production editor, who skillfully kept the process moving and us authors on track; Kerri Wilson, our photo researcher, who was so effective in finding photos to bring chemistry to life for students; and Roxy Wilson (University of Illinois), who so ably coordinated the difficult job of working out solutions to the end-of-chapter exercises. Finally, we wish to thank our families and friends for their love, support, encouragement, and patience as we brought this thirteenth edition to completion.

Theodore L. Brown Department of Chemistry University of Illinois at Urbana-Champaign Urbana, IL 61801

tlbrown@illinois.edu or tlbrown1@earthlink.net

H. Eugene LeMay, Jr. Department of Chemistry University of Nevada Reno, NV 89557

lemay@unr.edu

Bruce E. Bursten

Department of Chemistry University of Tennessee Knoxville, TN 37996 bbursten@utk.edu

Catherine J. Murphy Department of Chemistry University of Illinois at Urbana-Champaign Urbana, IL 61801

murphycj@illinois.edu.

Patrick M. Woodward Department of Chemistry and Biochemistry

The Ohio State University Columbus, OH 43210 woodward@chemistry. ohio-state.edu

Matthew W. Stoltzfus Department of Chemistry and Biochemistry

The Ohio State University Columbus, OH 43210 stoltzfus.5@osu.edu

Preface xxix

List of Resources

For Students

MasteringChemistry^®

(http://www.masteringchemistry.com)

^{MasteringChemistry}_® is the most effective, widely used online tutorial, homework and assessment system for chemistry. It helps instructors maximize class time with customizable, easy to-assign, and automatically graded assessments that motivate students to learn outside of class and arrive prepared for lecture. These assessments can easily be customized and personalized by instructors to suit their individual teaching style. The pow

erful gradebook provides unique insight into student and class performance even before the first test. As a result, instructors can spend class time where students need it most.

Pearson eText The integration of Pearson eText within ^{MasteringChemistry}_® gives students with eTexts easy access to the electronic text when they are logged into ^{MasteringChemistry}®^{. Pearson eText pages look exactly like} the printed text, offering powerful new functionality for students and instructors. Users can create notes, highlight text in different colors, create bookmarks, zoom, view in single-page or two-page view, and more.

Students Guide (0-321-94928-5) Prepared by James C. Hill of California State University. This book assists students through the text material with chapter overviews, learning objectives, a review of key terms, as well as self-tests with answers and explanations. This edition also features MCAT practice questions.

Solutions to Red Exercises (0-321-94926-9) Prepared by Roxy Wilson of the University of Illinois, Urbana-Champaign. Full solutions to all the red-numbered exercises in the text are provided. (Short answers to red exercises are found in the appendix of the text.)

Solutions to Black Exercises (0-321-94927-7) Prepared by Roxy Wilson of the University of Illinois, Urbana-Champaign. Full solutions to all the black-numbered exercises in the text are provided.

Laboratory Experiments (0-321-94991-9) Prepared by John H. Nelson of the University of Nevada, and Michael Lufaso of the University of North Florida with contributions by Matthew Stoltzfus of The Ohio State University. This manual contains 40 finely tuned experiments chosen to introduce students to basic lab techniques and to illustrate core chemical principles. This new edition has been revised with the addition of four brand new experiments to correlate more tightly with the text. You can also customize these labs through Catalyst, our custom database program. For more information, visit http://www. pearsoncustom.com/custom-library/

For Instructors

Solutions to Exercises (0-321-94925-0) Prepared by Roxy Wilson of the University of Illinois, Urbana-Champaign. This manual contains all end-of-chapter exercises in the text. With an instructor’s permission, this manual may be made available to students.

Online Instructor Resource Center (0-321-94923-4) This resource provides an integrated collection of resources to help instructors make efficient and effective use of their time. It features all artwork from the text, including figures and tables in PDF format for high-resolution printing, as well as five ^{prebuilt PowerPoint}_™ presentations. The first presentation contains the images embedded within PowerPoint slides. The second includes a complete lecture outline that is modifiable by the user. The final three presentations contain worked “in-chapter” sample exercises and questions to be used with Classroom Response Systems. The Instructor Resource Center also contains movies, animations, and electronic files of the Instructor Resource Manual, as well as the Test Item File.

TestGen Testbank (0-321-94924-2) Prepared by Andrea Leonard of the University of Louisiana. The Test Item File now provides a selection of more than 4,000 test questions with 200 new questions in the thirteenth edition and 200 additional algorithmic questions.

Online Instructor Resource Manual (0-321-94929-3) Prepared by Linda Brunauer of Santa Clara University and Elzbieta Cook of Louisiana State University. Organized by chapter, this manual offers detailed lecture outlines and complete descriptions of all available lecture demonstrations, interactive media assets, common student misconceptions, and more.

Annotated Instructor’s Edition to Laboratory Experiments (0-321-98608-3) Prepared by John H. Nelson of the University of Nevada, and Michael Lufaso of the University of North Florida with contributions by Matthew Stoltzfus of the Ohio State University. This AIE combines the full student lab manual with appendices covering the proper disposal of chemical waste, safety instructions for the lab, descriptions of standard lab equipment, answers to questions, and more.

WebCT Test Item File (IRC download only) 0-321-94931-5

Blackboard Test Item File (IRC download only) 0-321-94930-7

About the Authors

THE BROWN/LEMAY/BURSTEN/

MURpHY/WOODWARD/STOLTzfUS

AUTHOR TEAM values collaboration as an

integral component to overall success. While each author

brings unique talent, research interests, and teaching

experiences, the team works together to review and

develop the entire text. It is this collaboration that keeps

the content ahead of educational trends and contributes

to continuous innovations in teaching and learning

throughout the text and technology. Some of the new

key features in the thirteenth edition and accompanying

MasteringChemistry^® course are highlighted on the

following pages. 

THEODORE L. BROWN received his Ph.D. from Michigan State University in 1956. Since then, he has been a member of the faculty of the University of Illinois, Urbana-Champaign, where he is now Professor of Chemistry, Emeritus. He served as Vice Chancellor for Research, and Dean of The Graduate College, from 1980 to 1986, and as Founding Director of the Arnold and Mabel Beckman Institute for Advanced Science and Technology from 1987 to 1993. Professor Brown has been an Alfred P. Sloan Foundation Research Fellow and has been awarded a Guggenheim Fellowship. In 1972 he was awarded the American Chemical Society Award for Research in Inorganic Chemistry and received the American Chemical Society Award for Distinguished Service in the Advancement of Inorganic Chemistry in 1993. He has been elected a Fellow of the American Association for the Advancement of Science, the American Academy of Arts and Sciences, and the American Chemical Society.

H. EUGENE LEMAY, JR., received his B.S. degree in Chemistry from Pacific Lutheran University (Washington) and his Ph.D. in Chemistry in 1966 from the University of Illinois, Urbana-Champaign. He then joined the faculty of the University of Nevada, Reno, where he is currently Professor of Chemistry, Emeritus. He has enjoyed Visiting Professorships at the University of North Carolina at Chapel Hill, at the University College of Wales in Great Britain, and at the University of California, Los Angeles. Professor LeMay is a popular and effective teacher, who has taught thousands of students during more than 40 years of university teaching. Known for the clarity of his lectures and his sense of humor, he has received several teaching awards, including the University Distinguished Teacher of the Year Award (1991) and the first Regents’ Teaching Award given by the State of Nevada Board of Regents (1997).

BRUCE E. BURSTEN received his Ph.D. in Chemistry from the University of Wisconsin in 1978. After two years as a National Science Foundation Postdoctoral Fellow at Texas A&M University, he joined the faculty of The Ohio State University, where he rose to the rank of Distinguished University Professor. In 2005, he moved to the University of Tennessee, Knoxville, as Distinguished Professor of Chemistry and Dean of the College of Arts and Sciences. Professor Bursten has been a Camille and Henry Dreyfus Foundation Teacher-Scholar and an Alfred P. Sloan Foundation Research Fellow, and he is a Fellow of both the American Association for the Advancement of Science and the American Chemical Society. At Ohio State he has received the University Distinguished Teaching Award in 1982 and 1996, the Arts and Sciences Student Council Outstanding Teaching Award in 1984, and the University Distinguished Scholar Award in 1990. He received the Spiers Memorial Prize and Medal of the Royal Society of Chemistry in 2003, and the Morley Medal of the Cleveland Section of the American Chemical Society in 2005. He was President of the American Chemical Society for 2008. In addition to his teaching and service activities, Professor Bursten’s research program focuses on compounds of the transition-metal and actinide elements.

CATHERINE J. MURpHY received two B.S. degrees, one in Chemistry and one in Biochemistry, from the University of Illinois, Urbana-Champaign, in 1986. She received her Ph.D. in Chemistry from the University of Wisconsin in 1990. She was a National Science Foundation and National Institutes of Health Postdoctoral Fellow at the California Institute of Technology from 1990 to 1993. In 1993, she joined the faculty of the University of South Carolina, Columbia, becoming the Guy F. Lipscomb Professor of Chemistry in 2003. In 2009 she moved to the University of Illinois, Urbana-Champaign, as the Peter C. and Gretchen Miller Markunas Professor of Chemistry. Professor Murphy has been honored for both research and teaching as a Camille Dreyfus Teacher-Scholar, an Alfred P. Sloan Foundation Research Fellow, a Cottrell Scholar of the Research Corporation, a National Science Foundation CAREER Award winner, and a subsequent NSF Award for Special Creativity. She has also received a USC Mortar Board Excellence in Teaching Award, the USC Golden Key Faculty Award for Creative Integration of Research and Undergraduate Teaching, the USC Michael J. Mungo Undergraduate Teaching Award, and the USC Outstanding Undergraduate Research Mentor Award. Since 2006, Professor Murphy has served as a Senior Editor for the Journal of Physical Chemistry. In 2008 she was elected a Fellow of the American Association for the Advancement of Science. Professor Murphy’s research program focuses on the synthesis and optical properties of inorganic nanomaterials, and on the local structure and dynamics of the DNA double helix.

pATRICK M. WOODWARD received B.S. degrees in both Chemistry and Engineering from Idaho State University in 1991. He received a M.S. degree in Materials Science and a Ph.D. in Chemistry from Oregon State University in 1996. He spent two years as a postdoctoral researcher in the Department of Physics at Brookhaven National Laboratory. In 1998, he joined the faculty of the Chemistry Department at The Ohio State University where he currently holds the rank of Professor. He has enjoyed visiting professorships at the University of Bordeaux in France and the University of Sydney in Australia. Professor Woodward has been an Alfred P. Sloan Foundation Research Fellow and a National Science Foundation CAREER Award winner. He currently serves as an Associate Editor to the Journal of Solid State Chemistry and as the director of the Ohio REEL program, an NSF-funded center that works to bring authentic research experiments into the laboratories of first- and second-year chemistry classes in 15 colleges and universities across the state of Ohio. Professor Woodward’s research program focuses on understanding the links between bonding, structure, and properties of solid-state inorganic functional materials.

MATTHEW W. STOLTzfUS received his B.S. degree in Chemistry from Millersville University in 2002 and his Ph. D. in Chemistry in 2007 from The Ohio State University. He spent two years as a teaching postdoctoral assistant for the Ohio REEL program, an NSF-funded center that works to bring authentic research experiments into the general chemistry lab curriculum in 15 colleges and universities across the state of Ohio. In 2009, he joined the faculty of Ohio State where he currently holds the position of Chemistry Lecturer. In addition to lecturing general chemistry, Stoltzfus accepted the Faculty Fellow position for the Digital First Initiative, inspiring instructors to offer engaging digital learning content to students through emerging technology. Through this initiative, he developed an iTunes U general chemistry course, which has attracted over 120,000 students from all over the world. Stoltzfus has received several teaching awards, including the inaugural Ohio State University 2013 Provost’s Award for Distinguished Teaching by a Lecturer and he is recognized as an Apple Distinguished Educator.

Data-Driven Analytics A New Direction in Chemical Education A^{uthors traditionally revise roughly 25% of the end of chapter questions when producing}

a new edition. These changes typically involve modifying numerical variables/identities of chemical formulas to make them “new” to the next batch of students. While these changes are appropriate for the printed version of the text, one of the strengths of MasteringChemistry® is its ability to randomize variables so that every student receives a “different” problem. Hence, the effort

which authors have historically put into changing variables can now be used to improve questions. In order to make informed decisions, the author team consulted the massive reservoir of data available through MasteringChemistry® to revise their question bank. In particular, they analyized which problems were frequently assigned and why; they paid careful attention to the amount of time it took students to work through a problem (flagging those that took longer than expected) and they observed the wrong answer submissions and hints used (a measure used to calculate the difficulty of problems). This “metadata” served as a starting point for the discussion of which end of chapter questions should be changed.

For example, the breadth of ideas presented in Chapter 9 challenges students to understand three-dimensional visualization while simultaneously introducing several new concepts (particu larly VSEPR, hybrids, and Molecular Orbital theory) that challenge their critical thinking skills. In revising the exercises for the chapter, the authors drew on the metadata as well as their own experi ence in assigning Chapter 9 problems in Mastering Chemistry. From these analyses, we were able to articulate two general revision guidelines.

1. Improve coverage of topic areas that were underutilized: In Chapter 9, the authors noticed that there was a particularly low usage rate for questions concerning Molecular Orbital Theory. Based on the metadata and their own teaching experience with Mastering, they recognized an opportunity

to expand the coverage of MO theory. Two

brand new exercises that emphasize the basics of

MO theory were the result of this analysis

including the example below. This strategy

was replicated throughout the entire book.

2. Revise the least assigned existing problems. Much of the appeal of MasteringChemistry® for students is the immediate feedback they get when they hit submit, which also provides an opportunity to confront any misconceptions right away. For instructors, the appeal is that these problems are automatically graded. Essay questions fail to provide these advantages since they must be graded by an instructor before a student may receive feedback. Wherever possible, we revised current essay questions to include automatically graded material.

Bottom Line: The revision of the end of chapter questions in this edition is informed by robust data-driven analytics providing a new level of pedagogically-sound assessments for your students, all while making the time they spend working these problems even more valuable.

Helping Students Think Like Scientists

Design an Experiment

Starting with Chapter 3, every chapter will feature a Design an Experiment exercise. The goal

of these exercises is to challenge students to think like a scientist, imagining what kind of data

needs to be collected and what sort of experimental procedures will provide them the data

needed to answer the question. These exercises tend to be integrative, forcing students to draw

on many of the skills they have learned in the current and previous chapters.

Design an Experiment topics include:

Ch 3: Formation of Sulfur Oxides

Ch 4: Identification of Mysterious White Powders Ch 5: Joule Experiment

Ch 6: Photoelectric Effect and Electron Configurations Ch 7: Chemistry of Potassium Superoxide Ch 8: Benzene Resonance

Ch 9: Colors of Organic Dyes

Ch 10: Identification of an Unknown Noble Gas Ch 11: Hydraulic Fluids

Ch 12: Polymers

Ch 13: Volatile Solvent Molecules

Go figure

Go Figure questions encourage students to stop and analyze the artwork in the text, for conceptual understanding. “Voice Balloons” in selected figures help students break down and understand the components of the

image. These questions are also available in MasteringChemistry®. The number of

Go Figure questions in the thirteenth edition has increased by 25%.

Ch 14: Reaction Kinetics via Spectrophotometry

Ch 15: Beer’s Law and Visible-Light Spectroscopy Ch 16:  Acidity/Basicity of an Unknown Liquid

Ch 17: Understanding Differences in pKa

Ch 18: Effects of Fracking on Groundwater

Ch 19: Drug Candidates and the Equilibrium Constant Ch 20: Voltaic Cells

Ch 21: Discovery and Properties of Radium

Ch 22: Identification of Unknowns

Ch 23: Synthesis and Characterization of a Coordination Compound Ch 24: Quaternary Structure in Proteins

practice Exercises

A major new feature of this edition is the addition of a second Practice Exercise to accompany each Sample Exercise within the chapters. The new Practice Exercises are multiple-choice with correct answers provided for the students in an appendix. Specific wrong answer feedback, written by the authors, will be available in MasteringChemistry® The primary goal of the new Practice Exercise feature is to provide students with an additional problem to test mastery of the concepts in the text and to address the most common conceptual misunderstandings. To ensure the questions touched on the most common student misconceptions, the authors consulted the ACS Chemistry Concept inventory before writing their questions.

Give It Some Thought (GIST) questions

These informal, sharply-focused exercises allow students the opportunity to gauge whether they are “getting it” as they read the text. The number of GIST questions has increased throughout the text as well as in MasteringChemistry®.

Active and Visual

T^{he most effective learning happens when students actively participate and interact with}

material in order to truly internalize key concepts. The Brown/Lemay/Bursten/Murphy/ Woodward/Stoltzfus author team has spent decades refining their text based on educational research to the extent that it has largely defined how the general chemistry course is taught. With the thirteenth edition, these authors have extended this tradition by giving each student a way to personalize their learning experience through MasteringChemistry®. The MasteringChemistry® course for Brown/Lemay/Bursten/Murphy/Woodward/Stoltzfus evolves learning and technology usage far beyond the lecture-homework model. Many of these resources can be used pre-lecture, during class, and for assessment while providing each student with a personalized learning experi ence which gives them the greatest chance of succeeding.

Learning Catalytics

Learning Catalytics™ is a “bring your own device” student engagement, assessment, and classroom intelligence system. With Learning Catalytics™ you can:

• Assess students in real time, using open-ended tasks to probe student understanding. • Understand immediately where students are and adjust your lecture accordingly. • Improve your students’ critical-thinking skills.

• Access rich analytics to understand student performance.

• Add your own questions to make Learning Catalytics™ fit your course exactly. • Manage student interactions with intelligent grouping and timing.

Learning Catalytics™ is a technology that has grown out of twenty years of cutting-edge research, innovation, and implementation of interactive teaching and peer instruction.

Learning Catalytics™ will be included with the purchase of MasteringChemistry® with eText.

pause and predict Videos

Author Dr. Matt Stoltzfus created Pause and Predict Videos. These videos engage students by prompting them to submit a prediction about the outcome of an experiment or demonstration before seeing the final result. A set of assignable tutorials, based on these videos, challenge students to transfer their understanding of the demonstration to related scenarios. These videos are also available in web- and mobile-friendly formats through the study area of MasteringChemistry® and in the Pearson eText.

NEW! Simulations, assignable in MasteringChemistry®, include those developed by the PhET Chemistry Group, and the leading authors in simulation development covering some of the most difficult chemistry concepts.

Adaptive

MasteringChemistry® has always been personalized and adaptive on a question level by providing error-specific feedback based on actual student responses; however, Mastering now includes two

new adaptive assignment types—Adaptive Follow-Up Assignments and Dynamic Study Modules.

Adaptive follow-Up Assignments

Instructors have the ability to assign adaptive follow-up assignments. Content delivered to

students as part of adaptive learning will be automatically personalized for each individual

based on strengths and weaknesses identified by his or her performance on Mastering

parent assignments.

Question sets in the Adaptive

Follow-Up Assignments continu

ously adapt to each student’s needs,

making efficient use of study time.

Dynamic Study Modules

NEW! Dynamic Study Modules, designed to enable students to study effectively on their own as well as help students quickly access and learn the nomenclature they need to be successful in chemistry.

These modules can be accessed on smartphones, tablets, and computers and results can be tracked in the MasteringChemistry® Gradebook. Here’s how it works:

1. Students receive an initial set of questions and benefit from the metacognition involved with asking them to indicate how confident they are with their answer. 2. After answering each set of questions, students review their answers. 3. Each question has explanation material that reinforces the correct answer response and addresses the misconceptions found in the wrong answer choices.

4. Once students review the explanations, they are presented with a new set of questions. Students cycle through this dynamic process of test-learn-retest until they achieve mastery of the material.

1

Introduction: Matter

and Measurement

In the title of this book we refer to chemistry as the central science. This title reflects the fact that much of what goes on in the world around us involves chemistry. The changes that produce the brilliant colors of tree leaves in the fall, the electrical energy that powers a cell phone, the spoilage of foods left standing at room temperature, and the many ways in which our bodies use the foods we consume are all everyday examples of chemical processes.

Chemistry is the study of matter and the changes that matter undergoes. As you progress

in your study, you will come to see how chemical principles operate in all aspects of our

lives, from everyday activities like food preparation to more complex processes such as

those that operate in the environment. We use chemical principles to understand a host of

phenomena, from the role of salt in our diet to the workings of a lithium ion battery.

This first chapter provides an overview of what chemistry is about and what chem

ists do. The “What’s Ahead” list gives an overview of the chapter organization and of

some of the ideas we will consider.

▶ THE BEAUTIFUL COLORS that develop

1.1 ^| The Study of Chemistry

Chemistry is at the heart of many changes we see in the world around us, and it ac counts for the myriad of different properties we see in matter. To understand how these changes and properties arise, we need to look far beneath the surfaces of our everyday observations.

What’s

in trees in the fall appear when the tree ceases to produce chlorophyll, which imparts the green color to the leaves during the summer. Some of the color we see has been in the leaf all summer, and some develops from the action of sunlight on the leaf as the chlorophyll disappears.

Ahead

1.1 The Study of Chemistry We begin with a brief description of what chemistry is, what chemists do, and why it is useful to learn chemistry.

1.2 Classifications of Matter Next, we examine some fundamental ways to classify matter, distinguishing between

1.3 Properties of Matter We then consider different characteristics, or properties, used to characterize, identify, and separate substances, distinguishing between chemical and physical properties.

1.4 Units of Measurement We observe that many properties rely on quantitative measurements involving numbers and units. The units of measurement used throughout science are those of the metric system.

pure substances and mixtures and between elements and compounds.

1.5 Uncertainty in Measurement We observe that the uncertainty inherent in all measured quantities is expressed by the number of significant figures used to report the quantity. Significant figures are also used to express the uncertainty associated with calculations involving measured quantities.

1.6 Dimensional Analysis We recognize that units as well as numbers are carried through calculations and that obtaining correct units for the result of a calculation is an important way to check whether the calculation is correct.

4 chapter 1 Introduction: Matter and Measurement

The Atomic and Molecular Perspective of Chemistry

Chemistry is the study of the properties and behavior of matter. Matter is the physical

material of the universe; it is anything that has mass and occupies space. A property is

any characteristic that allows us to recognize a particular type of matter and to distinguish

it from other types. This book, your body, the air you are breathing, and the clothes you

are wearing are all samples of matter. We observe a tremendous variety of matter in our

world, but countless experiments have shown that all matter is comprised of combina

tions of only about 100 substances called elements. One of our major goals will be to relate

the properties of matter to its composition, that is, to the particular elements it contains.

Chemistry also provides a background for understanding the properties of matter

in terms of atoms, the almost infinitesimally small building blocks of matter. Each ele

ment is composed of a unique kind of atom. We will see that the properties of matter re

late to both the kinds of atoms the matter contains (composition) and the arrangements

of these atoms (structure).

In molecules, two or more atoms are joined in specific shapes. Throughout this text

you will see molecules represented using colored spheres to show how the atoms are con

nected (▼ Figure 1.1). The color provides a convenient way to distinguish between atoms

of different elements. For example, notice that the molecules of ethanol and ethylene gly

col in Figure 1.1 have different compositions and structures. Ethanol contains one oxygen

atom, depicted by one red sphere. In contrast, ethylene glycol contains two oxygen atoms.

Even apparently minor differences in the composition or structure of molecules

can cause profound differences in properties. For example, let’s compare ethanol and

ethylene glycol, which appear in Figure 1.1 to be quite similar. Ethanol is the alcohol in

beverages such as beer and wine, whereas ethylene glycol is a viscous liquid used as au

tomobile antifreeze. The properties of these two substances differ in many ways, as do

their biological activities. Ethanol is consumed throughout the world, but you should

never consume ethylene glycol because it is highly toxic. One of the challenges chemists

undertake is to alter the composition or structure of molecules in a controlled way, cre

ating new substances with different properties. For example, the common drug aspirin,

shown in Figure 1.1, was first synthesized in 1897 in a successful attempt to improve on

a natural product extracted from willow bark that had long been used to alleviate pain.

Every change in the observable world—from boiling water to the changes that occur

as our bodies combat invading viruses—has its basis in the world of atoms and molecules.

Go Figure

Which of the molecules in the figure has the most carbon atoms? How many are there in that molecule?= H = O = C

Oxygen

Water

Ethanol

Carbon dioxide Ethylene glycol

Aspirin

▲ Figure 1.1 Molecular models. The white, black, and red spheres represent atoms of hydrogen, carbon, and oxygen, respectively.

section 1.1 The Study of Chemistry 5

Thus, as we proceed with our study of chemistry, we will find ourselves thinking in two

realms: the macroscopic realm of ordinary-sized objects 1macro = large2 and the submi

croscopic realm of atoms and molecules. We make our observations in the macroscopic

world, but to understand that world, we must visualize how atoms and molecules behave at

the submicroscopic level. Chemistry is the science that seeks to understand the properties

and behavior of matter by studying the properties and behavior of atoms and molecules.

Give It Some Thought

(a) Approximately how many elements are there?

(b) What submicroscopic particles are the building blocks of matter?

Why Study Chemistry?

Chemistry lies near the heart of many matters of public concern, such as improvement

of health care, conservation of natural resources, protection of the environment, and the

supply of energy needed to keep society running. Using chemistry, we have discovered

and continually improved upon pharmaceuticals, fertilizers and pesticides, plastics, solar

panels, LEDs, and building materials. We have also discovered that some chemicals are

potentially harmful to our health or the environment. This means that we must be sure

that the materials with which we come into contact are safe. As a citizen and consumer,

it is in your best interest to understand the effects, both positive and negative, that chem

icals can have, and to arrive at a balanced outlook regarding their uses.

You may be studying chemistry because it is an essential part of your curriculum.

Your major might be chemistry, or it could be biology, engineering, pharmacy, agricul

ture, geology, or some other field. Chemistry is central to a fundamental understand

ing of governing principles in many science-related fields. For example, our interactions

with the material world raise basic questions about the materials around us. ▼ Figure 1.2

illustrates how chemistry is central to several different realms of modern life.

Energy

Solar panels are composed

Biochemistry

The ash of the re y results

of specially treated silicon.

Technology

LED’s (light emitting diodes) are formed from elements such as gallium, arsenic and phosphorus.

from a chemical reaction in the insect.

Chemistry

Medicine

Connectors and tubing for medical procedures such as intravenous injections are made from plastics highly resistant to chemical attack.

▲ Figure 1.2 Chemistry is central to our understanding of the world around us.

6 chapter 1 Introduction: Matter and Measurement

Chemistry Put to Work

Chemistry and the Chemical

Industry

Chemistry is all around us. Many people are familiar with household chemicals, particularly kitchen chemicals such as those shown in ▶ Figure 1.3. However, few realize the size and importance of the chemical industry. Worldwide sales of chemicals and related prod

ucts manufactured in the United States total approximately $585 bil lion annually. Sales of pharmaceuticals total another $180 billion. The chemical industry employs more than 10% of all scientists and engi

neers and is a major contributor to the U.S. economy. Vast amounts of industrial chemicals are produced each year. ▼ Table 1.1 lists several of the chemicals produced in highest vol umes in the United States. Notice that they all serve as raw materi als for a variety of uses, including the manufacture and processing of metals, plastics, fertilizers, and other goods.

Who are chemists, and what do they do? People who have degrees in chemistry hold a variety of positions in industry, govern ment, and academia. Those in industry work as laboratory chem ists, developing new products (research and development); analyzing materials (quality control); or assisting customers in using products (sales and service). Those with more experience or training may work as managers or company directors. Chemists are important members of the scientific workforce in government (the National Institutes of Health, Department of Energy, and Environmental Protection Agency all employ chemists) and at universities. A chem istry degree is also good preparation for careers in teaching, medi cine, biomedical research, information science, environmental work, technical sales, government regulatory agencies, and patent law.

Fundamentally, chemists do three things: (1) make new types of matter: materials, substances, or combinations of substances with

desired properties; (2) measure the properties of matter; and (3) develop models that explain and/or predict the properties of matter. One chem ist, for example, may work in the laboratory to discover new drugs. An other may concentrate on the development of new instrumentation to measure properties of matter at the atomic level. Other chemists may use existing materials and methods to understand how pollutants are transported in the environment or how drugs are processed in the body. Yet another chemist will develop theory, write computer code, and run computer simulations to understand how molecules move and react. The collective chemical enterprise is a rich mix of all of these activities.

▲ Figure 1.3 Common chemicals employed in home food production.

Table 1.1 Several of the Top Chemicals Produced by the U.S. Chemical Industry* Annual Production

Chemical Formula

(Billions of Pounds) Principal End Uses

Sulfuric acid H₂SO₄ 70 Fertilizers, chemical manufacturing Ethylene C₂H₄ 50 Plastics, antifreeze

Lime CaO 45 Paper, cement, steel

Propylene C₃H₆ 35 Plastics

Ammonia NH₃ 18 Fertilizers

Chlorine Cl₂ 21 Bleaches, plastics, water purification Phosphoric acid H₃PO₄ 20 Fertilizers

Sodium hydroxide NaOH 16 Aluminum production, soap

1.2 ^| Classifications of Matter

Let’s begin our study of chemistry by examining two fundamental ways in which mat ter is classified. Matter is typically characterized by (1) its physical state (gas, liquid, or solid) and (2) its composition (whether it is an element, a compound, or a mixture).

*Data from Chemical & Engineering News, July 2, 2007, pp. 57, 60, American Chemical Society; data online from U.S. Geological Survey.

section 1.2 Classifications of Matter 7

States of Matter

A sample of matter can be a gas, a liquid, or a solid. These

three forms, called the states of matter, differ in some of their observable properties. A gas (also known as vapor) has no fixed volume or shape; rather, it uniformly fills its

Go Figure

In which form of water are the water molecules farthest apart?

container. A gas can be compressed to occupy a smaller volume, or it can expand to occupy a larger one. A liq uid has a distinct volume independent of its container, and assumes the shape of the portion of the container it occupies. A solid has both a definite shape and a definite volume. Neither liquids nor solids can be compressed to any appreciable extent.

The properties of the states of matter can be under stood on the molecular level (▶ Figure 1.4). In a gas the molecules are far apart and moving at high speeds, col liding repeatedly with one another and with the walls of the container. Compressing a gas decreases the amount of space between molecules and increases the frequency of collisions between molecules but does not alter the size or shape of the molecules. In a liquid, the molecules are packed closely together but still move rapidly. The rapid movement allows the molecules to slide over one an other; thus, a liquid pours easily. In a solid the molecules are held tightly together, usually in definite arrangements in which the molecules can wiggle only slightly in their otherwise fixed positions. Thus, the distances between molecules are similar in the liquid and solid states, but the two states differ in how free the molecules are to move

Water vapor

Ice

Liquid water

around. Changes in temperature and/or pressure can lead to conversion from one state of matter to another, illus trated by such familiar processes as ice melting or water vapor condensing.

Pure Substances

▲ Figure 1.4 The three physical states of water—water vapor, liquid water, and ice. We see the liquid and solid states but cannot see the gas (vapor) state. The red arrows show that the three states of matter interconvert.

Most forms of matter we encounter—the air we breathe (a gas), the gasoline we burn in our cars (a liquid), and the sidewalk we walk on (a solid)—are not chemically pure. We can, however, separate these forms of matter into pure substances. A pure substance (usually referred to simply as a substance) is matter that has distinct properties and a composition that does not vary from sample to sample. Water and table salt (sodium chloride) are examples of pure substances.

All substances are either elements or compounds. Elements are substances that cannot be decomposed into simpler substances. On the molecular level, each element is composed of only one kind of atom [Figure 1.5(a and b)]. Compounds are substances composed of two or more elements; they contain two or more kinds of atoms [Figure 1.5(c)]. Water, for example, is a compound composed of two elements: hydrogen and oxygen.

Figure 1.5(d) shows a mixture of substances. Mixtures are combinations of two or more substances in which each substance retains its chemical identity.

Elements

Currently, 118 elements are known, though they vary widely in abundance. Hydrogen constitutes about 74% of the mass in the Milky Way galaxy, and helium constitutes 24%. Closer to home, only five elements—oxygen, silicon, aluminum, iron, and calcium—account for over 90% of Earth’s crust (including oceans and atmosphere), and only three—oxygen, carbon, and hydrogen—account for over 90% of the mass of the human body (Figure 1.6).

8 chapter 1 Introduction: Matter and Measurement

Go Figure

How do the molecules of a compound differ from the molecules of an element?

(a) Atoms of an element (d) Mixture of elements

(b) Molecules of an element

(c) Molecules

of a compound

and a compound

Only one kind of atom is in any element. Compounds must have at

least two kinds of atoms.

▲ Figure 1.5 Molecular comparison of elements, compounds, and mixtures.

▼ Table 1.2 lists some common elements, along with the chemical symbols

Go Figure

Name two significant differences between the elemental composition of Earth’s crust and the elemental composition of the human body.

used to denote them. The symbol for each element consists of one or two letters, with the first letter capitalized. These symbols are derived mostly from the Eng lish names of the elements, but sometimes they are derived from a foreign name instead (last column in Table 1.2). You will need to know these symbols and learn others as we encounter them in the text.

All of the known elements and their symbols are listed on the front

Calcium 3.4%

Iron 4.7%

Aluminum 7.5%

Other 9.2%

Silicon

inside cover of this text in a table known as the periodic table. In the periodic table the elements are arranged in columns so that closely related elements are grouped together. We describe the periodic table in more detail in Section 2.5 and consider the periodically repeating properties of the elements in Chapter 7.

Oxygen

49.5%

Earth’s crust

25.7%

Compounds

Most elements can interact with other elements to form compounds. For example, when hydrogen gas burns in oxygen gas, the elements hydrogen and oxygen combine to form the compound water. Conversely, water can be decom posed into its elements by passing an electrical current through it (▶ Figure 1.7).

Oxygen 65%

Other

^7%Hydrogen 10%

Carbon

18%

Table 1.2 Some Common Elements and Their Symbols Carbon C Aluminum Al Copper Cu (from cuprum) Fluorine F Bromine Br Iron Fe (from ferrum) Hydrogen H Calcium Ca Lead Pb (from plumbum) Iodine I Chlorine Cl Mercury Hg (from hydrargyrum) Nitrogen N Helium He Potassium K (from kalium)

Human body

▲ Figure 1.6 Relative abundances of elements.* Elements in percent by mass in Earth’s crust (including oceans and atmosphere) and the human body.

Oxygen O Lithium Li Silver Ag (from argentum) Phosphorus P Magnesium Mg Sodium Na (from natrium) Sulfur S Silicon Si Tin Sn (from stannum)

*U.S. Geological Survey Circular 285, U.S Department of the Interior.

section 1.2 Classifications of Matter 9

Go Figure

How are the relative gas volumes collected in the two tubes related to the relative number of gas molecules in the tubes?

Oxygen gas, O₂

Water, H₂O Hydrogen gas, H₂

▲ Figure 1.7 Electrolysis of water. Water decomposes into its component elements, hydrogen

and oxygen, when an electrical current is passed through it. The volume of hydrogen, collected

in the right test tube, is twice the volume of oxygen.

Pure water, regardless of its source, consists of 11% hydrogen and 89% oxygen by mass.

This macroscopic composition corresponds to the molecular composition, which

consists of two hydrogen atoms combined with one oxygen atom:

Hydrogen atom (written H)

Oxygen atom (written O)

Water molecule (written H₂O)

The elements hydrogen and oxygen themselves exist naturally as diatomic (two atom) molecules:

Oxygen molecule Hydrogen molecule

(written O₂) (written H₂)

As seen in ▼ Table 1.3, the properties of water bear no resemblance to the proper ties of its component elements. Hydrogen, oxygen, and water are each a unique sub stance, a consequence of the uniqueness of their respective molecules.

Table 1.3 Comparison of Water, Hydrogen, and Oxygen

Water Hydrogen Oxygen

State^a Liquid Gas Gas

Normal boiling point 100 °C -253 °C -183 °C

Density^a 1000 g/L 0.084 g/L 1.33 g/L

Flammable No Yes No

^aAt room temperature and atmospheric pressure.

10 chapter 1 Introduction: Matter and Measurement

The observation that the elemental composition of a compound is always the same

is known as the law of constant composition (or the law of definite proportions).

French chemist Joseph Louis Proust (1754–1826) first stated the law in about 1800.

Although this law has been known for 200 years, the belief persists among some peo

ple that a fundamental difference exists between compounds prepared in the labora

tory and the corresponding compounds found in nature. However, a pure compound

has the same composition and properties under the same conditions regardless of its

source. Both chemists and nature must use the same elements and operate under the

same natural laws. When two materials differ in composition or properties, either they

are composed of different compounds or they differ in purity.

Give It Some Thought

Hydrogen, oxygen, and water are all composed of molecules. What is it about a

molecule of water that makes it a compound, whereas hydrogen and oxygen are

elements?

Mixtures

Most of the matter we encounter consists of mixtures of different substances. Each sub

stance in a mixture retains its chemical identity and properties. In contrast to a pure

substance, which by definition has a fixed composition, the composition of a mixture

can vary. A cup of sweetened coffee, for example, can contain either a little sugar or a

lot. The substances making up a mixture are called components of the mixture.

Some mixtures do not have the same composition, properties, and appearance

throughout. Rocks and wood, for example, vary in texture and appearance in any

typical sample. Such mixtures are heterogeneous [▼ Figure 1.8(a)]. Mixtures that are

uniform throughout are homogeneous. Air is a homogeneous mixture of nitrogen,

oxygen, and smaller amounts of other gases. The nitrogen in air has all the proper

ties of pure nitrogen because both the pure substance and the mixture contain the

same nitrogen molecules. Salt, sugar, and many other substances dissolve in water to

form homogeneous mixtures [Figure 1.8(b)]. Homogeneous mixtures are also called

solutions. Although the term solution conjures an image of a liquid, solutions can be

solids, liquids, or gases.

▶ Figure 1.9 summarizes the classification of matter into elements, compounds,

and mixtures.

(a) (b)

▲ Figure 1.8 Mixtures. (a) Many common materials, including rocks, are heterogeneous mixtures.

This photograph of granite shows a heterogeneous mixture of silicon dioxide and other metal

oxides. (b) Homogeneous mixtures are called solutions. Many substances, including the blue solid

shown here [copper(II) sulfate], dissolve in water to form solutions.

section 1.3 Properties of Matter 11

Matter

NO YES

Is it uniform

throughout?

Heterogeneous

_mixture Homogeneous

Does it have a

NO YES

variable

composition?

_{Pure substance} Homogeneous

mixture

(solution)

Does it contain

NO YES

more than one

kind of atom?

Element Compound

▲ Figure 1.9 Classification of matter. All pure matter is classified ultimately as either an element

or a compound.

Sample

Exercise 1.1 Distinguishing among Elements, Compounds, and Mixtures

“White gold” contains gold and a “white” metal, such as palladium. Two samples of white gold

differ in the relative amounts of gold and palladium they contain. Both samples are uniform in

composition throughout. Use Figure 1.9 to classify white gold.

Solution

Because the material is uniform throughout, it is homogeneous. Because its composition differs for the two samples, it cannot be a compound. Instead, it must be a homogeneous mixture.

Practice Exercise 1

Which of the following is the correct description of a cube of material cut from the inside of an apple?

(a) It is a pure compound.

(b) It consists of a homogenous mixture of compounds. 1.3 ^| Properties of Matter

(c) It consists of a heterogeneous mixture of compounds. (d) It consists of a heterogeneous mixture of elements and compounds.

(e) It consists of a single compound in different states.

Practice Exercise 2

Aspirin is composed of 60.0% carbon, 4.5% hydrogen, and 35.5% oxygen by mass, regardless of its source. Use Figure 1.9 to classify aspirin.

Every substance has unique properties. For example, the properties listed in Table 1.3 allow us to distinguish hydrogen, oxygen, and water from one another. The properties of matter can be categorized as physical or chemical. Physical properties can be ob served without changing the identity and composition of the substance. These proper ties include color, odor, density, melting point, boiling point, and hardness. Chemical properties describe the way a substance may change, or react, to form other substances. A common chemical property is flammability, the ability of a substance to burn in the presence of oxygen.

Some properties, such as temperature and melting point, are intensive properties. Intensive properties do not depend on the amount of sample being examined and are particularly useful in chemistry because many intensive properties can be used to identify substances. Extensive properties depend on the amount of sample, with two examples being mass and volume. Extensive properties relate to the amount of substance present.

12 chapter 1 Introduction: Matter and Measurement

Give It Some Thought

When we say that lead is a denser metal than aluminum, are we talking about an

extensive or intensive property?

Physical and Chemical Changes

The changes substances undergo are either physical or chemical. During a physical

change, a substance changes its physical appearance but not its composition. (That is, it

is the same substance before and after the change.) The evaporation of water is a physi

cal change. When water evaporates, it changes from the liquid state to the gas state, but

it is still composed of water molecules, as depicted in Figure 1.4. All changes of state

(for example, from liquid to gas or from liquid to solid) are physical changes.

In a chemical change (also called a chemical reaction), a substance is transformed

into a chemically different substance. When hydrogen burns in air, for example, it under

goes a chemical change because it combines with oxygen to form water (▼ Figure 1.10).

H₂ O₂

Burn

H₂ O₂ H₂O

▲ Figure 1.10 A chemical reaction.

Chemical changes can be dramatic. In the account that follows, Ira Remsen, author

of a popular chemistry text published in 1901, describes his first experiences with

chemical reactions. The chemical reaction that he observed is shown in ▼ Figure 1.11.

▲ Figure 1.11 The chemical reaction between a copper penny and nitric acid. The dissolved copper produces the blue-green solution; the reddish brown gas produced is nitrogen dioxide.

section 1.3 Properties of Matter 13

While reading a textbook of chemistry, I came upon the statement “nitric acid acts upon

copper,” and I determined to see what this meant. Having located some nitric acid, I had

only to learn what the words “act upon” meant. In the interest of knowledge I was even

willing to sacrifice one of the few copper cents then in my possession. I put one of them

on the table, opened a bottle labeled “nitric acid,” poured some of the liquid on the cop

per, and prepared to make an observation. But what was this wonderful thing which I

beheld? The cent was already changed, and it was no small change either. A greenish-blue

liquid foamed and fumed over the cent and over the table. The air became colored dark

red. How could I stop this? I tried by picking the cent up and throwing it out the window.

I learned another fact: nitric acid acts upon fingers. The pain led to another unpremedi

tated experiment. I drew my fingers across my trousers and discovered nitric acid acts

upon trousers. That was the most impressive experiment I have ever performed. I tell of it

even now with interest. It was a revelation to me. Plainly the only way to learn about such

remarkable kinds of action is to see the results, to experiment, to work in the laboratory.*

Give It Some Thought

Which of these changes are physical and which are chemical? Explain.

(a) Plants make sugar from carbon dioxide and water.

(b) Water vapor in the air forms frost.

(c) A goldsmith melts a nugget of gold and pulls it into a wire.

Separation of Mixtures

We can separate a mixture into its components by taking advantage of differences in

their properties. For example, a heterogeneous mixture of iron filings and gold filings

could be sorted by color into iron and gold. A less tedious approach would be to use a

magnet to attract the iron filings, leaving the gold ones behind. We can also take ad

vantage of an important chemical difference between these two metals: Many acids dis

solve iron but not gold. Thus, if we put our mixture into an appropriate acid, the acid

would dissolve the iron and the solid gold would be left behind. The two could then be

separated by filtration (▶ Figure 1.12). We would have to use other chemical reactions,

which we will learn about later, to transform the dissolved iron back into metal.

An important method of separating the components of a homogeneous mixture

is distillation, a process that depends on the different abilities of substances to form

gases. For example, if we boil a solution of salt and water, the water evaporates, forming

a gas, and the salt is left behind. The gaseous water can be converted back to a liquid on

the walls of a condenser, as shown in ▼ Figure 1.13.

2

Boiling the solution

1

vaporizes the water

Water is condensed,

and then collected in

the receiving ask

Condenser

Salt water

Cold water

out

Cold water

in

After water has boiled away,

3

pure sodium chloride remains

Pure water

in receiving ask

▲ Figure 1.12 Separation by filtration. A mixture of a solid and a liquid is poured

▲ Figure 1.13 Distillation. Apparatus for separating a sodium chloride solution (salt water) into its components.

*Remsen, Ira, The Principles of Theoretical Chemistry, 1887.

through filter paper. The liquid passes through the paper while the solid remains on the paper.

14 chapter 1 Introduction: Matter and Measurement

Go Figure

Is the separation of a, b, and c in Figure 1.14 a physical or chemical process? I II III

^{ow of solven}t

Solvent

Mixture of

compounds

(a + b + c)

Adsorbent

(stationary

phase)

Glass

a a

b + c

b

c

Compounds a, b, and c ^wool Stopcock

are adsorbed to different

degrees on the solid

stationary phase

▲ Figure 1.14 Separation of three substances using column chromatography.

The differing abilities of substances to adhere to the surfaces of solids can also

be used to separate mixtures. This ability is the basis of chromatography, a technique

shown in ▲ Figure 1.14.

1.4 ^| Units of Measurement

Many properties of matter are quantitative, that is, associated with numbers. When a

number represents a measured quantity, the units of that quantity must be specified.

To say that the length of a pencil is 17.5 is meaningless. Expressing the number with its

units, 17.5 centimeters (cm), properly specifies the length. The units used for scientific

▲ Figure 1.15 Metric units. Metric measurements are increasingly common in the United States, as exemplified by the volume printed on this soda can in both English units (fluid ounces, fl oz) and metric units (milliliters, mL).

A Closer Look

measurements are those of the metric system.

The metric system, developed in France during the late eighteenth century, is used as the system of measurement in most countries. The United States has traditionally used the English system, although use of the metric system has become more common (◀ Figure 1.15).

The Scientific Method

Where does scientific knowledge come from? How is it acquired? How do we know it is reliable? How do scientists add to it, or modify it? There is nothing mysterious about how scientists work. The first idea to keep in mind is that scientific knowledge is gained through observations of the natural world. A principal aim of the scientist is to organize these observations, by identifying patterns and regularity, making measurements, and associating one set of observations with another. The next step is to ask why nature behaves in the manner we observe. To answer this question, the scientist constructs a model,

known as a hypothesis, to explain the observations. Initially the hy pothesis is likely to be pretty tentative. There could be more than one reasonable hypothesis. If a hypothesis is correct, then certain results and observations should follow from it. In this way hypotheses can stimulate the design of experiments to learn more about the system being studied. Scientific creativity comes into play in thinking of hy potheses that are fruitful in suggesting good experiments to do, ones that will shed new light on the nature of the system.

As more information is gathered, the initial hypotheses get winnowed down. Eventually just one may stand out as most consis tent with a body of accumulated evidence. We then begin to call this

hypothesis a theory, a model that has predictive powers, and that ac counts for all the available observations. A theory also generally is consistent with other, perhaps larger and more general theories. For example, a theory of what goes on inside a volcano has to be consistent with more general theories regarding heat transfer, chemistry at high temperature, and so forth.

We will be encountering many theories as we proceed through this book. Some of them have been found over and over again to be consistent with observations. However, no theory can be proven to be absolutely true. We can treat it as though it is, but there always remains a possibility that there is some respect in which a theory is wrong. A famous example is Einstein’s theory of relativ ity. Isaac Newton’s theory of mechanics yielded such precise results for the mechanical behavior of matter that no exceptions to it were found before the twentieth century. But Albert Einstein showed that Newton’s theory of the nature of space and time is incorrect. Einstein’s theory of relativity represented a fundamental shift in how we think of space and time. He predicted where the exceptions to predictions based on Newton’s theory might be found. Although only small departures from Newton’s theory were predicted, they were observed. Einstein’s theory of relativity became accepted as the correct model. However, for most uses, Newton’s laws of motion are quite accurate enough.

The overall process we have just considered, illustrated in ▶ Figure 1.16, is often referred to as the scientific method. But there is no single scientific method. Many factors play a role in advancing scientific knowledge. The one unvarying requirement is that our explanations be consistent with observations, and that they depend solely on natural phenomena.

When nature behaves in a certain way over and over again, under all sorts of different conditions, we can summarize that behavior in a scientific law. For example, it has been repeatedly observed that in a chemical reaction there is no change in the total mass of the materials reacting as compared with the materi als that are formed; we call this observation the Law of Conserva

tion of Mass. It is important to make a distinction between a theory and a scientific law. The latter simply is a statement of what always

SI Units

section 1.4 Units of Measurement 15

happens, to the best of our knowledge. A theory, on the other hand, is an explanation for what happens. If we discover some law fails to hold true, then we must assume the theory underlying that law is wrong in some way.

Related Exercises: 1.60, 1.82

Collect information via

observations of natural

phenomena and experiments

Formulate one or more

explanatory hypotheses

Perform experiments to

test the hypotheses

Use the most successful

hypotheses to formulate

a theory

Repeatedly test theory.

Modify as needed to match

experimental results, or reject.

▲ Figure 1.16 The scientific method.

In 1960 an international agreement was reached specifying a particular choice of metric units for use in scientific measurements. These preferred units are called SI units, after the French Système International d’Unités. This system has seven base units from which all other units are derived (▼ Table 1.4). In this chapter we will consider the base units for length, mass, and temperature.

Table 1.4 SI Base Units

Physical Quantity Name of Unit Abbreviation Mass Kilogram kg Length Meter m

Time Second s or sec Temperature Kelvin K

Amount of substance Mole mol Electric current Ampere A or amp Luminous intensity Candela cd

16 chapter 1 Introduction: Matter and Measurement

Give It Some Thought

The package of a fluorescent bulb for a table lamp lists the light output in terms

of lumens, lm. Which of the seven SI units would you expect to be part of the

definition of a lumen?

With SI units, prefixes are used to indicate decimal fractions or multiples of vari

ous units. For example, the prefix milli- represents a 10^-3 fraction, one-thousandth, of

a unit: A milligram (mg) is 10^-3 gram (g), a millimeter (mm) is 10^-3 meter (m), and so

forth. ▼ Table 1.5 presents the prefixes commonly encountered in chemistry. In using

SI units and in working problems throughout this text, you must be comfortable using

exponential notation. If you are unfamiliar with exponential notation or want to review

it, refer to Appendix A.1.

Although non–SI units are being phased out, some are still commonly used by sci

entists. Whenever we first encounter a non–SI unit in the text, the SI unit will also be

given. The relations between the non–SI and SI units we will use most frequently in this

text appear on the back inside cover. We will discuss how to convert from one to the

other in Section 1.6.

Table 1.5 Prefixes Used in the Metric System and with SI Units

Prefix Abbreviation Meaning Example

^{Peta P} 10^{15 1 petawatt (PW)} = 1 * 10¹⁵ watts^a

^{Tera T} 10^{12 1 terawatt (TW)} = 1 * 10¹² watts

^{Giga G} 10^{9 1 gigawatt (GW)} = 1 * 10⁹ watts

^{Mega M} 10^{6 1 megawatt (MW)} = 1 * 10⁶ watts

^{Kilo k} 10^{3 1 kilowatt (kW)} = 1 * 10³ watts

^{Deci d} 10^{-1 1 deciwatt (dW)} = 1 * 10^-1 watt

^{Centi c} 10^{-2 1 centiwatt (cW)} = 1 * 10^-2 watt

^{Milli m} 10^{-3 1 milliwatt (mW)} = 1 * 10^-3 watt

^Micro m^b 10^-6 1 microwatt 1mW2 = 1 * 10^-6 watt

^{Nano n} 10^{-9 1 nanowatt (nW)} = 1 * 10^-9 watt

^{Pico p} 10^{-12 1 picowatt (pW)} = 1 * 10^-12 watt

^{Femto f} 10^{-15 1 femtowatt (fW)} = 1 * 10^-15 watt

^{Atto a} 10^{-18 1 attowatt (aW)} = 1 * 10^-18 watt

^{Zepto z} 10^{-21 1 zeptowatt (zW)} = 1 * 10^-21 watt

^aThe watt (W) is the SI unit of power, which is the rate at which energy is either generated

or consumed. The SI unit of energy is the joule (J); _{1 J = 1 kg} # _m²>s² and 1 W = 1 J>s.

^bGreek letter mu, pronounced “mew.”

Give It Some Thought How many mg are there in 1 mg?

Length and Mass

section 1.4 Units of Measurement 17

The SI base unit of length is the meter, a distance slightly longer than a yard. Mass* is a measure of the amount of material in an object. The SI base unit of mass is the kilogram (kg), which is equal to about 2.2 pounds (lb). This base unit is unusual because it uses a pre fix, kilo-, instead of the word gram alone. We obtain other units for mass by adding prefixes to the word gram.

Sample

exercise 1.2 Using SI Prefixes

What is the name of the unit that equals (a) 10^-9 gram, (b) 10^-6 second, (c) 10^-3 meter?

Solution

We can find the prefix related to each power of ten in Table 1.5: (a) nanogram, ng; (b) microsec ond, ms; (c) millimeter, mm.

Practice Exercise 1

Which of the following weights would you expect to be suitable for weighing on an ordinary bathroom scale?

(a) 2.0 * 10⁷ mg, (b) 2500 mg, (c) 5 * 10^-4 kg, (d) 4 * 10⁶ cg, (e) 5.5 * 10⁸ dg.

Practice Exercise 2

(a) How many picometers are there in 1 m? (b) Express 6.0 * 10³ m using a prefix to replace the power of ten. (c) Use exponential notation to express 4.22 mg in grams. (d) Use decimal notation to express 4.22 mg in grams.

Temperature

Temperature, a measure of the hotness or coldness of an object, is a physical property that determines the direction of heat flow. Heat always flows spontaneously from a sub stance at higher temperature to one at lower temperature. Thus, the influx of heat we feel when we touch a hot object tells us that the object is at a higher temperature than our hand.

The temperature scales commonly employed in science are the Celsius and Kelvin scales. The Celsius scale was originally based on the assignment of 0 °C to the freezing point of water and 100 °C to its boiling point at sea level (Figure 1.17).

*Mass and weight are often incorrectly thought to be the same. The weight of an object is the force that is exerted on its mass by gravity. In space, where gravitational forces are very weak, an astronaut can be weightless, but he or she cannot be massless. The astronaut’s mass in space is the same as it is on Earth.

18 chapter 1 Introduction: Matter and Measurement

Go Figure

True or false: The “size” of a degree on the Celsius scale is the same as the “size” of a degree on the Kelvin scale.

373 K

_{100 degree-interval}s

100 °C

_{100 degree-interval}s

212 °F 98.6 °F

Water boils

_{180 degree-interval}s

310 K 37.0 °C

Normal body temperature

273 K

0 °C

32 °F

Water freezes

Kelvin scale Celsius scale

Fahrenheit scale

▲ Figure 1.17 Comparison of the Kelvin, Celsius, and Fahrenheit temperature scales.

The Kelvin scale is the SI temperature scale, and the SI unit of temperature is the

kelvin (K). Zero on the Kelvin scale is the lowest attainable temperature, referred to

as absolute zero. On the Celsius scale, absolute zero has the value, -273.15 °C. The

Celsius and Kelvin scales have equal-sized units—that is, a kelvin is the same size as a

degree Celsius. Thus, the Kelvin and Celsius scales are related according to

K = °C + 273.15 [1.1]

The freezing point of water, 0 °C, is 273.15 K (Figure 1.17). Notice that we do not use a

degree sign 1°2 with temperatures on the Kelvin scale.

The common temperature scale in the United States is the Fahrenheit scale, which

is not generally used in science. Water freezes at 32 °F and boils at 212 °F. The Fahren

heit and Celsius scales are related according to

_{°C =} ⁵_{9 1°F - 322 or °F =} ⁹₅ 1°C2 + 32 [1.2]

Sample

exercise 1.3 Converting Units of Temperature

A weather forecaster predicts the temperature will reach 31 °C. What is this temperature (a) in K,

(b) in °F?

Solution

(a) Using Equation 1.1, we have K = 31 + 273 = 304 K. (b) Using Equation 1.2, we have

_{°F =} ⁹₅1312 + 32 = 56 + 32 = 88 °F.

Practice Exercise 1

Using Wolfram Alpha (http://www.wolframalpha.com/) or some other reference, determine which of these elements would be

liquid at 525 K (assume samples are protected from air): (a) bismuth, Bi; (b) platinum, Pt; (c) selenium, Se; (d) calcium, Ca; (e) copper, Cu.

Practice Exercise 2

Ethylene glycol, the major ingredient in antifreeze, freezes at -11.5 °C. What is the freezing point in (a) K, (b) °F?

section 1.4 Units of Measurement 19

Derived SI Units

The SI base units are used to formulate derived units. A derived unit is obtained by

multiplication or division of one or more of the base units. We begin with the defin ing equation for a quantity and, then substitute the appropriate base units. For exam ple, speed is defined as the ratio of distance traveled to elapsed time. Thus, the SI unit for speed—m/s, read “meters per second”—is a derived unit, the SI unit for distance (length), m, divided by the SI unit for time, s. Two common derived units in chemistry are those for volume and density.

Volume

The volume of a cube is its length cubed, length³. Thus, the derived SI unit of volume is the SI unit of length, m, raised to the third power. The cubic meter, m³, is the volume of a cube that is 1 m on each edge (▶ Figure 1.18). Smaller units, such as cubic cen timeters, cm³ (sometimes written cc), are frequently used in chemistry. Another vol ume unit used in chemistry is the liter (L), which equals a cubic decimeter, dm³, and is slightly larger than a quart. (The liter is the first metric unit we have encountered that is not an SI unit.) There are 1000 milliliters (mL) in a liter, and 1 mL is the same volume as 1 cm³: 1 mL = 1 cm³. The devices used most frequently in chemistry to measure vol ume are illustrated in ▼ Figure 1.19.

Syringes, burettes, and pipettes deliver amounts of liquids with more precision than graduated cylinders. Volumetric flasks are used to contain specific volumes of liquid.

Give It Some Thought

Which of the following quantities represents volume measurement: 15 m²; 2.5 * 10² m³; 5.77 L>s? How do you know?

Density

Density is defined as the amount of mass in a unit volume of a substance: _{Density =} mass

_volume[1.3]

Go Figure

How many 1-L bottles are required to contain 1 m³ of liquid?

_{1 m} 1 m

1 m

1 dm³ = 1L

1 cm³ = 1mL

1 cm

1 cm 1 cm

▲ Figure 1.18 Volume relationships. The volume occupied by a cube 1 m on each edge is one cubic meter, 1 m³. Each cubic meter contains 1000 dm³. One liter is the same volume as one cubic decimeter, 1 L = 1 dm³. Each cubic decimeter contains 1000 cubic centimeters, 1 dm³ = 1000 cm³. One cubic centimeter equals one milliliter, 1 cm³ = 1 mL.

mL 100 90

80

70

60

50

40

30

20

10

These deliver variable volumes Pipette delivers a speci c volume

mL 0

1

2

3

4

5

45

46

47

48

49

50

Stopcock,

a valve to

control the

liquid ow

Volumetric ask contains a speci c volume

Graduated cylinder

Syringe

Burette

Pipette Volumetric ask

▲ Figure 1.19 Common volumetric glassware.

20 chapter 1 Introduction: Matter and Measurement

The densities of solids and liquids are commonly expressed in either grams per

Table 1.6 Densities of Selected Substances at 25 °C

^Substance Density ^1g,^cm32 Air 0.001

Balsa wood 0.16

Ethanol 0.79

Water 1.00

Ethylene glycol 1.09

Table sugar 1.59

Table salt 2.16

Iron 7.9

Gold 19.32

Sample

cubic centimeter 1g>cm³2 or grams per milliliter 1g>mL2. The densities of some com mon substances are listed in ◀ Table 1.6. It is no coincidence that the density of water is 1.00 g>mL; the gram was originally defined as the mass of 1 mL of water at a specific temperature. Because most substances change volume when they are heated or cooled, densities are temperature dependent, and so temperature should be specified when re porting densities. If no temperature is reported, we assume 25 °C, close to normal room temperature.

The terms density and weight are sometimes confused. A person who says that iron weighs more than air generally means that iron has a higher density than air—1 kg of air has the same mass as 1 kg of iron, but the iron occupies a smaller volume, thereby giving it a higher density. If we combine two liquids that do not mix, the less dense liq

uid will float on the denser liquid.

Exercise 1.4 Determining Density and Using Density to Determine Volume or Mass (a) Calculate the density of mercury if 1.00 * 10² g occupies a volume of 7.36 cm³.

(b) Calculate the volume of 65.0 g of liquid methanol (wood alcohol) if its density is 0.791 g>mL. (c) What is the mass in grams of a cube of gold 1density = 19.32 g>cm³2 if the length of the cube is 2.00 cm?

Solution

(a) We are given mass and volume, so Equation 1.3 yields _{volume =} 1.00 * 10² g

Practice Exercise 1

Platinum, Pt, is one of the rarest of the metals. Worldwide annual

_{Density =} mass

7.36 cm³ = 13.6 g>cm³

production is only about 130 tons. (a) Platinum has a density of 21.4 g>cm³. If thieves were to steal platinum from a bank using a

(b) Solving Equation 1.3 for volume and then using the given mass _{density =} 65.0 g

small truck with a maximum payload of 900 lb, how many 1 L bars of the metal could they make off with? (a) 19 bars, (b) 2 bars,

_{and density gives Volume =} mass

_{0.791 g>mL} = 82.2 mL

(c) 42 bars, (d) 1 bar, (e) 47 bars.

(c) We can calculate the mass from the volume of the cube and its density. The volume of a cube is given by its length cubed:

Volume = 12.00 cm2³ = 12.002³ cm³ = 8.00 cm³

Solving Equation 1.3 for mass and substituting the volume and density of the cube, we have

Mass = volume * density = 18.00 cm³2119.32 g>cm³2 = 155 g

Chemistry Put to Work

Chemistry in the News

Because chemistry is so central to our lives, reports on matters of chem ical significance appear in the news nearly every day. Some reports tell of breakthroughs in the development of new pharmaceuticals, materi als, and processes. Others deal with energy, environmental, and public safety issues. As you study chemistry, you will develop the skills to better understand the importance of chemistry in your life. Here are summa

ries of a few recent stories in which chemistry plays an important role. Clean energy from fuel cells. In fuel cells, the energy of a chemical reaction is converted directly into electrical energy. Although fuel cells have long been known as potentially valuable sources of electrical energy, their costs have kept them from widespread use. However, recent advanc es in technology have brought fuel cells to the fore as sources of reliable and clean electrical power in certain critical situations. They are

Practice Exercise 2

(a) Calculate the density of a 374.5-g sample of copper if it has a volume of 41.8 cm³. (b) A student needs 15.0 g of ethanol for an experiment. If the density of ethanol is 0.789 g>mL, how many milliliters of ethanol are needed? (c) What is the mass, in grams, of 25.0 mL of mercury 1density = 13.6 g>mL2?

especially valuable in powering data centers which consume large amounts of electrical power that must be absolutely reliable. For example, failure of electrical power at a major data center for a company such as Amazon, eBay, or Apple could be calamitous for the company and its customers.

eBay recently contracted to build the next phase of its major data cen ter in Utah, utilizing solid–state fuel cells as the source of electrical power. The fuel cells, manufactured by Bloom Energy, a Silicon Valley startup, are large industrial devices about the size of a refrigerator (▶ Figure 1.20). The eBay installation utilizes biogas, which consists of methane and other fuel gases derived from landfills and farms. The fuel is combined with oxygen, and the mixture run through a special solid–state device to pro duce electricity. Because the electricity is being produced close to the data center, transmission of the electrical power from source to consumption is more efficient. In contrast to electrical backup systems employed in the past, the new power source will be the primary source of power, operating

▲ Figure 1.20 Solid-State fuel cells manufactured by Bloom Energy.

24 hours per day, every day of the year. The eBay facility in Utah is the largest nonelectric utility fuel cell installation in the nation. It generates 6 megawatts of power, enough to power about 6000 homes.

Regulation of greenhouse gases. In 2009 the Environmental Pro tection Agency (EPA) took the position that, under the provisions of the Clean Air Act, it should regulate emissions of “greenhouse” gases. Greenhouse gases are substances that have the potential to alter the global climate because of their ability to trap long–wavelength radia tion at Earth’s surface. (This subject is covered in detail in Section 18.2.) Greenhouse gases include carbon dioxide 1CO₂2, methane 1CH₄2, and nitrous oxide 1N₂O2, as well as other substances. The EPA decision was challenged in the courts by several states, industry organizations, and conservative groups. In a major victory for the EPA, the federal court of appeals of the District of Columbia in July 2012 upheld the agen cy’s position. This case is interesting in part because of the grounds on which the EPA policy was challenged, and the way the court responded. The plaintiffs argued that the EPA improperly based its decision on as sessments from the Intergovernmental Panel on Climate Change, the U.S. Global Climate Change program, and reports from the Nation al Research Council, rather than on citing the findings of individual research programs in the published literature. The court replied that “it makes no difference that much of the scientific evidence in large part consisted of ‘syntheses’ of individual studies and research. This is how science works. EPA is not required to re-prove the existence of the atom every time it approaches a scientific question.”*

This is an important example of an interaction between science and social policy in our complex, modern society. When other than purely scientific interests are involved, questions about science’s reli ability and objectivity are bound to arise.

Anesthesia. In the period around the 1840s it became recognized that certain substances, notably ether, chloroform, and nitrous oxide, could induce a state in which the patient had no awareness of bodily pain. You can imagine how joyfully these new discoveries were received by people who had to undergo surgery that would otherwise be unbear

*U.S. Court of Appeals for the District of Columbia , Case No. 09-1322.

section 1.4 Units of Measurement 21

ably painful. The word anesthesia was sug

gested by Oliver Wendell Holmes, Sr. in 1846

to describe the state in which a person lacks

awareness, either total or of a particular part of

the body. In time chemists were able to iden

tify certain organic compounds that produced

anesthesia without being severely toxic.

More than 40 million patients in North

America each year undergo medical proce

dures that call for anesthesia. The anesthet

ics used today are most often injected into

the blood stream rather than inhaled as a gas.

Several organic substances have been identi

fied as effective anesthetics. While modern

anesthetics are generally quite safe, they must

be administered with care, because they can

affect breathing, blood pressure, and heart

function. Every drug has a therapeutic index,

the ratio of the smallest dose that would be fa

tal to the smallest dose that gives the desired

therapeutic effect. Naturally, one wants the

therapeutic index for any drug to be as large as

possible. Anesthetics have generally low thera

peutic indices, which means that they must

be administered carefully and with constant

monitoring. The death of the entertainer Mi

chael Jackson in June 2009 from an overdose of propofol, a widely used anesthetic (▼ Figure 1.21), illustrates how dangerous such drugs can be when not properly administered. Propofol very quickly renders a pa tient unconscious and affects breathing. Hence its use must be carefully monitored by a person trained in anesthesiology.

Despite a great deal of research, it is still not clear how anesthet ics actually work. It is a near-universal characteristic of life that spe cies ranging from tadpoles to humans can be reversibly immobilized. The search for the mechanisms by which this occurs is important, be cause it may lead us not only to safer anesthetics, but also to deeper understanding of what we mean by consciousness itself.

▲ Figure 1.21 Propofol, an anesthetic.

22 chapter 1 Introduction: Matter and Measurement

Go Figure

How would the darts be positioned on the target for the case of “good accuracy, poor precision”?

Good accuracy

Good precision

Poor accuracy

Good precision

Poor accuracy

Poor precision

▲ Figure 1.22 Precision and accuracy. High precision can be achieved on a scale like this one, which has 0.1 milligram accuracy.

1.5 ^| Uncertainty in Measurement

Two kinds of numbers are encountered in scientific work: exact numbers (those whose values are known exactly) and inexact numbers (those whose values have some uncer tainty). Most of the exact numbers we will encounter in this book have defined values. For example, there are exactly 12 eggs in a dozen, exactly 1000 g in a kilogram, and ex actly 2.54 cm in an inch. The number 1 in any conversion factor, such as 1 m = 100 cm or 1 kg = 2.2046 lb, is an exact number. Exact numbers can also result from counting objects. For example, we can count the exact number of marbles in a jar or the exact number of people in a classroom.

Numbers obtained by measurement are always inexact. The equipment used to measure quantities always has inherent limitations (equipment errors), and there are differences in how different people make the same measurement (human errors). Sup pose ten students with ten balances are to determine the mass of the same dime. The ten measurements will probably vary slightly for various reasons. The balances might be calibrated slightly differently, and there might be differences in how each student reads the mass from the balance. Remember: Uncertainties always exist in measured quantities.

Give It Some Thought

Which of the following is an inexact quantity?

(a) the number of people in your chemistry class

(b) the mass of a penny

(c) the number of grams in a kilogram

Precision and Accuracy

The terms precision and accuracy are often used in discussing the uncertainties of mea sured values. Precision is a measure of how closely individual measurements agree with one another. Accuracy refers to how closely individual measurements agree with the correct, or “true,” value. The dart analogy in ◀ Figure 1.22 illustrates the difference between these two concepts.

In the laboratory we often perform several “trials” of an experiment and aver age the results. The precision of the measurements is often expressed in terms of the standard deviation (Appendix A.5), which reflects how much the individual measurements differ from the average. We gain confidence in our measurements if we obtain nearly the same value each time—that is, when the standard deviation is small. Figure 1.22 reminds us, however, that precise measurements can be inac curate. For example, if a very sensitive balance is poorly calibrated, the masses we measure will be consistently either high or low. They will be inaccurate even if they are precise.

Significant Figures

Suppose you determine the mass of a dime on a balance capable of measuring to the nearest 0.0001 g. You could report the mass as 2.2405 { 0.0001 g. The { no tation (read “plus or minus”) expresses the magnitude of the uncertainty of your measurement. In much scientific work we drop the { notation with the under standing that there is always some uncertainty in the last digit reported for any mea sured quantity.

▶ Figure 1.23 shows a thermometer with its liquid column between two scale marks. We can read the certain digits from the scale and estimate the uncertain one. Seeing that the liquid is between the 25° and 30 °C marks, we estimate the temperature to be 27 °C, being uncertain of the second digit of our measurement. By uncertain we mean that the temperature is reliably 27 °C and not 28° or 26 °C, but we can’t say that it is exactly 27 °C.

100 °C 80 °C 60 °C 40 °C 20 °C 0 °C

30 °C

27 °C 25 °C

section 1.5 Uncertainty in Measurement 23

◀ Figure 1.23 Uncertainty and significant

figures in a measurement.

Second digit in 27 °C is

estimated and therefore

uncertain

All digits of a measured quantity, including the uncertain one, are called signifi cant figures. A measured mass reported as 2.2 g has two significant figures, whereas one reported as 2.2405 g has five significant figures. The greater the number of signifi cant figures, the greater the precision implied for the measurement.

Sample

Exercise 1.5 Relating Significant Figures to the Uncertainty of a Measurement What difference exists between the measured values 4.0 and 4.00 g?

Solution

The value 4.0 has two significant figures, whereas 4.00 has three. This difference implies that 4.0 has more uncertainty. A mass reported as 4.0 g indicates that the uncertainty is in the first decimal place. Thus, the mass is closer to 4.0 than to 3.9 or 4.1 g. We can rep resent this uncertainty by writing the mass as 4.0 { 0.1 g. A mass reported as 4.00 g indicates that the uncertainty is in the second decimal place. In this case the mass is closer to 4.00 than 3.99 or 4.01 g, and we can represent it as 4.00 { 0.01 g. (Without further information, we cannot be sure whether the difference in uncertain ties of the two measurements reflects the precision or the accuracy of the measurement.)

Give It Some Thought

Practice Exercise 1

Mo Farah won the 10,000 meter race in the 2012 Olympics with an official time of 27 minutes, 30.42 s. To the correct number of significant figures, what was Farah’s average speed in m/sec? (a) 0. 6059 m/s, (b) 1.65042 m/s, (c) 6.059064 m/s, (d) 0.165042 m/s, (e) 6.626192 m/s.

Practice Exercise 2

A sample that has a mass of about 25 g is weighed on a balance that has a precision of {0.001 g. How many significant figures should be reported for this measurement?

A digital bathroom scale gives you the following four readings in a row: 155.2, 154.8, 154.9, 154.8 lbs. How would you record your weight?

To determine the number of significant figures in a reported measurement, read the number from left to right, counting the digits starting with the first digit that is not zero. In any measurement that is properly reported, all nonzero digits are significant. Because zeros can be used either as part of the measured value or merely to locate the decimal point, they may or may not be significant:

1. Zeros between nonzero digits are always significant—1005 kg (four significant figures); 7.03 cm (three significant figures).

2. Zeros at the beginning of a number are never significant; they merely indicate the position of the decimal point—0.02 g (one significant figure); 0.0026 cm (two sig nificant figures).

24 chapter 1 Introduction: Matter and Measurement

3. Zeros at the end of a number are significant if the number contains a decimal

point—0.0200 g (three significant figures); 3.0 cm (two significant figures).

A problem arises when a number ends with zeros but contains no decimal point.

In such cases, it is normally assumed that the zeros are not significant. Exponential

notation (Appendix A.1) can be used to indicate whether end zeros are significant. For

example, a mass of 10,300 g can be written to show three, four, or five significant fig

ures depending on how the measurement is obtained:

1.03 * 10⁴ g (three significant figures)

1.030 * 10⁴ g (four significant figures)

1.0300 * 10⁴ g (five significant figures)

In these numbers all the zeros to the right of the decimal point are significant (rules

1 and 3). (The exponential term 10⁴ does not add to the number of significant

figures.)

Sample

Exercise 1.6 Assigning Appropriate Significant Figures

The state of Colorado is listed in a road atlas as having a population of 4,301,261 and an area of

104,091 square miles. Do the numbers of significant figures in these two quantities seem reason

able? If not, what seems to be wrong with them?

Solution

The population of Colorado must vary from day to day as people move in or out, are born, or die. Thus, the reported number suggests a much higher degree of accuracy than is possible. Secondly, it would not be feasible to actually count every individual resident in the state at any given time. Thus, the reported number suggests far greater precision than is possible. A reported number of 4,300,000 would better reflect the actual state of knowledge.

The area of Colorado does not normally vary from time to time, so the question here is whether the accuracy of the measurements is good to six significant figures. It would be possible to achieve such accuracy using satellite technology, provided the legal boundaries are known with sufficient accuracy.

Sample

Practice Exercise 1

Which of the following numbers in your personal life are exact numbers?

(a) Your cell phone number, (b) your weight, (c) your IQ, (d) your driver’s license number, (e) the distance you walked yesterday.

Practice Exercise 2

The back inside cover of the book tells us that there are 5280 ft in 1 mile. Does this make the mile an exact distance?

Exercise 1.7 Determining the Number of Significant Figures in a Measurement

How many significant figures are in each of the following numbers (assume that each number is a measured quantity)? (a) 4.003, (b) 6.023 * 10²³, (c) 5000.

Solution

(a) Four; the zeros are significant figures. (b) Four; the exponential term does not add to the number of significant figures. (c) One; we assume that the zeros are not significant when there is no decimal point shown. If the number has more significant figures, a decimal point should be employed or the number written in exponential notation. Thus, 5000. has four significant figures, whereas 5.00 * 10³ has three.

Practice Exercise 1

Sylvia feels as though she may have a fever. Her normal body temperature is 98.7 °F. She measures her body temperature with a

thermometer placed under her tongue and gets a value of 102.8 °F. How many significant figures are in this measurement? (a) Three, the number of degrees to the left of the decimal point; (b) four, the number of digits in the measured reading; (c) two, the number of digits in the difference between her current reading and her normal body temperature; (d) three, the number of digits in her normal body temperature; (e) one, the number of digits to the right of the decimal point in the measured value.

Practice Exercise 2

How many significant figures are in each of the following mea surements? (a) 3.549 g, (b) 2.3 * 10⁴ cm, (c) 0.00134 m³.

section 1.5 Uncertainty in Measurement 25

Significant Figures in Calculations

When carrying measured quantities through calculations, the least certain measure

ment limits the certainty of the calculated quantity and thereby determines the number

of significant figures in the final answer. The final answer should be reported with only

one uncertain digit. To keep track of significant figures in calculations, we will make

frequent use of two rules: one for addition and subtraction, and another for multiplica

tion and division.

1. For addition and subtraction, the result has the same number of decimal places as

the measurement with the fewest decimal places. When the result contains more

than the correct number of significant figures, it must be rounded off. Consider the

following example in which the uncertain digits appear in color:

This number limits 20.42 — two decimal places

the number of significant 1.322 — three decimal places

figures in the result ¡ 83.1 — one decimal place

104.842 — round off to one decimal place (104.8)

We report the result as 104.8 because 83.1 has only one decimal place.

2. For multiplication and division, the result contains the same number of sig

nificant figures as the measurement with the fewest significant figures. When

the result contains more than the correct number of significant figures, it must

be rounded off. For example, the area of a rectangle whose measured edge lengths

are 6.221 and 5.2 cm should be reported with two significant figures, 32 cm², even

though a calculator shows the product to have more digits:

Area = 16.221 cm215.2 cm2 = 32.3492 cm² 1 round off to 32 cm²

because 5.2 has two significant figures.

Notice that for addition and subtraction, decimal places are counted in determining how

many digits to report in an answer, whereas for multiplication and division, significant

figures are counted in determining how many digits to report.

In determining the final answer for a calculated quantity, exact numbers are as

sumed to have an infinite number of significant figures. Thus, when we say, “There are

12 inches in 1 foot,” the number 12 is exact, and we need not worry about the number

of significant figures in it.

In rounding off numbers, look at the leftmost digit to be removed:

• If the leftmost digit removed is less than 5, the preceding number is left unchanged.

Thus, rounding off 7.248 to two significant figures gives 7.2.

• If the leftmost digit removed is 5 or greater, the preceding number is increased by 1.

Rounding off 4.735 to three significant figures gives 4.74, and rounding 2.376 to

two significant figures gives 2.4.*

Give It Some Thought

A rectangular garden plot is measured to be 25.8 m by 18 m. Which of these

dimensions needs to be measured to greater accuracy to provide a more accurate

estimate of the area of the plot?

*Your instructor may want you to use a slight variation on the rule when the leftmost digit to be removed is

exactly 5, with no following digits or only zeros following. One common practice is to round up to the next

higher number if that number will be even and down to the next lower number otherwise. Thus, 4.7350

would be rounded to 4.74, and 4.7450 would also be rounded to 4.74.

26 chapter 1 Introduction: Matter and Measurement

Sample

Exercise 1.8 Determining the Number of Significant Figures in a Calculated Quantity

The width, length, and height of a small box are 15.5, 27.3, and 5.4 cm, respectively. Calculate the

volume of the box, using the correct number of significant figures in your answer.

Solution

In reporting the volume, we can show only as many significant figures as given in the dimension with the fewest significant figures, which is that for the height (two significant figures):

Volume = width * length * height

= 115.5 cm2127.3 cm215.4 cm2

= 2285.01 cm³ 1 2.3 * 10³ cm³

A calculator used for this calculation shows 2285.01, which we must round off to two significant figures. Because the resulting number is 2300, it is best reported in exponential notation, 2.3 * 10³, to clearly indicate two significant figures.

Sample

Practice Exercise 1

Ellen recently purchased a new hybrid car and wants to check her gas mileage. At an odometer setting of 651.1 mi, she fills the tank. At 1314.4 mi she requires 16.1 gal to refill the tank. Assuming that the tank is filled to the same level both times, how is the gas mile

age best expressed? (a) 40 mi/gal, (b) 41 mi/gal, (c) 41.2 mi/gal, (d) 41.20 mi/gal.

Practice Exercise 2

It takes 10.5 s for a sprinter to run 100.00 m. Calculate her average speed in meters per second and express the result to the correct number of significant figures.

Exercise 1.9 Determining the Number of Significant Figures in a Calculated Quantity

A vessel containing a gas at 25 °C is weighed, emptied, and then reweighed as depicted in

▼ Figure 1.24. From the data provided, calculate the density of the gas at 25 °C.

Solution

To calculate the density, we must know both the mass and the volume of the gas. The mass of the gas is just the difference in the masses of the full and empty container:

1837.63 - 836.252 g = 1.38 g

In subtracting numbers, we determine the number of significant fig

case each quantity has two decimal places. Thus, the mass of the gas, 1.38 g, has two decimal places.

Using the volume given in the question, 1.05 * 10³ cm³, and the defi nition of density, we have

_{volume =} 1.38 g

ures in our result by counting decimal places in each quantity. In this

_{Density =} mass

1.05 * 10³ cm³

Pump out gas

Volume: 1.05 × 10³ cm³

Mass: 837.63 g

Mass: 836.25 g

= 1.31 * 10^-3 g>cm³ = 0.00131 g>cm³

In dividing numbers, we determine the number of significant fig ures our result should contain by counting the number of significant figures in each quantity. There are three significant figures in our answer, corresponding to the number of significant figures in the two numbers that form the ratio. Notice that in this example, following the rules for determining significant figures gives an answer containing only three significant figures, even though the measured masses con tain five significant figures.

Practice Exercise 1

Which of the following numbers is correctly rounded to three significant figures, as shown in brackets? (a) 12,556 [12,500], (b) 4.5671 * 10^-9 34.567 * 10^-94, (c) 3.00072 [3.001], (d) 0.006739 [0.00674], (e) 5.4589 * 10⁵ 35.459 * 10⁵4.

Practice Exercise 2

If the mass of the container in the sample exercise (Figure 1.24) were measured to three decimal places before and after pumping

▲ Figure 1.24 Uncertainty and significant figures in a measurement.

out the gas, could the density of the gas then be calculated to four significant figures?

When a calculation involves two or more steps and you write answers for intermedi ate steps, retain at least one nonsignificant digit for the intermediate answers. This pro cedure ensures that small errors from rounding at each step do not combine to affect the final result. When using a calculator, you may enter the numbers one after another,

section 1.6 Dimensional Analysis 27

rounding only the final answer. Accumulated rounding-off errors may account for

small differences among results you obtain and answers given in the text for numerical

problems.

1.6 ^| Dimensional Analysis

Because measured quantities have units associated with them, it is important to keep

track of units as well as numerical values when using the quantities in calculations.

Throughout the text we use dimensional analysis in solving problems. In dimen

sional analysis, units are multiplied together or divided into each other along with

the numerical values. Equivalent units cancel each other. Using dimensional analysis

helps ensure that solutions to problems yield the proper units. Moreover, it provides

a systematic way of solving many numerical problems and of checking solutions for

possible errors.

The key to using dimensional analysis is the correct use of conversion factors to

change one unit into another. A conversion factor is a fraction whose numerator and

denominator are the same quantity expressed in different units. For example, 2.54 cm

and 1 in. are the same length: 2.54 cm = 1 in. This relationship allows us to write two

conversion factors:

2.54 cm

_{1 in. and} 1 in.

2.54 cm

We use the first factor to convert inches to centimeters. For example, the length in

centimeters of an object that is 8.50 in. long is

Desired unit

^{2.54 cm} Number of centimeters = (8.50 in.) = 21.6 cm

1 in.

Given unit

The unit inches in the denominator of the conversion factor cancels the unit

inches in the given data (8.50 inches), so that the centimeters unit in the numera

tor of the conversion factor becomes the unit of the final answer. Because the

numerator and denominator of a conversion factor are equal, multiplying any

quantity by a conversion factor is equivalent to multiplying by the number 1 and

so does not change the intrinsic value of the quantity. The length 8.50 in. is the

same as the length 21.6 cm.

In general, we begin any conversion by examining the units of the given data

and the units we desire. We then ask ourselves what conversion factors we have

available to take us from the units of the given quantity to those of the desired one.

When we multiply a quantity by a conversion factor, the units multiply and divide

as follows:

_{Given unit *}desired unit

_{given unit} = desired unit

If the desired units are not obtained in a calculation, an error must have been made

somewhere. Careful inspection of units often reveals the source of the error.

Sample

Exercise 1.10 Converting Units

If a woman has a mass of 115 lb, what is her mass in grams? (Use the relationships between units

given on the back inside cover of the text.)

Solution

Because we want to change from pounds to grams, we look for a relationship between these units

of mass. The conversion factor table found on the back inside cover tells us that 1 lb = 453.6 g.

28 chapter 1 Introduction: Matter and Measurement

To cancel pounds and leave grams, we write the conversion factor with grams in the numerator

Given: lb

_Use 453.6 g

1 lb

Find: g

and pounds in the denominator:

Mass in grams = 1115 lb_2a453.6 g

_{1 lb} b = 5.22 * 10⁴ g

The answer can be given to only three significant figures, the number of significant figures in 115 lb. The process we have used is diagrammed in the margin.

Practice Exercise 1

At a particular instant in time the Earth is judged to be 92,955,000 miles from the Sun. What is the distance in kilometers to four significant figures? (See back inside cover for conversion factor). (a) 5763 * 10⁴ km, (b) 1.496 * 10⁸ km, (c) 1.49596 * 10⁸ km, (d) 1.483 * 10⁴ km, (e) 57,759,000 km.

Practice Exercise 2

By using a conversion factor from the back inside cover, determine the length in kilometers of a 500.0-mi automobile race.

Strategies in Chemistry

Estimating Answers

Calculators are wonderful devices; they enable you to get to the wrong answer very quickly. Of course, that’s not the destination you want. You can take certain steps to avoid putting that wrong answer into your homework set or on an exam. One is to keep track of the units in a calculation and use the correct conversion factors. Second, you can do a quick mental check to be sure that your an swer is reasonable: you can try to make a “ballpark” estimate.

A ballpark estimate involves making a rough calculation using numbers that are rounded off in such a way that the arithmetic can be

done without a calculator. Even though this approach does not give an exact answer, it gives one that is roughly the correct size. By using di mensional analysis and by estimating answers, you can readily check the reasonableness of your calculations.

You can get better at making estimates by practicing in every day life. How far is it from your dorm room to the chemistry lecture hall? How much do your parents pay for gasoline per year? How many bikes are there on campus? If you respond “I have no idea” to these questions, you’re giving up too easily. Try estimating familiar quanti ties and you’ll get better at making estimates in science and in other aspects of your life where a misjudgment can be costly.

Give It Some Thought

How do we determine how many digits to use in conversion factors, such as the one between pounds and grams in Sample Exercise 1.10?

Using Two or More Conversion Factors

It is often necessary to use several conversion factors in solving a problem. As an ex ample, let’s convert the length of an 8.00-m rod to inches. The table on the back inside cover does not give the relationship between meters and inches. It does, however, give the relationship between centimeters and inches 11 in. = 2.54 cm2. From our knowl edge of SI prefixes, we know that 1 cm = 10^-2 m. Thus, we can convert step by step, first from meters to centimeters and then from centimeters to inches:

Given:

Use Use

Find:

m

1 cm

10⁻² m

^cm 1 in. 2.54 cm

in.

Combining the given quantity (8.00 m) and the two conversion factors, we have 10^-2 _{m b a} 1 in.

_{Number of inches = 18.00 m2a} 1 cm

_{2.54 cm} b = 315 in.

The first conversion factor is used to cancel meters and convert the length to centime ters. Thus, meters are written in the denominator and centimeters in the numerator.

section 1.6 Dimensional Analysis 29

The second conversion factor is used to cancel centimeters and convert the length to

inches, so it has centimeters in the denominator and inches, the desired unit, in the

numerator.

Note that you could have used 100 cm = 1 m as a conversion factor as well in the

second parentheses. As long as you keep track of your given units and cancel them

properly to obtain the desired units, you are likely to be successful in your calculations.

Sample

Exercise 1.11 Converting Units Using Two or More Conversion Factors

The average speed of a nitrogen molecule in air at 25 °C is 515 m>s. Convert this speed to miles per hour.

Solution

To go from the given units, m/s, to the desired units, mi/hr, we must convert meters to miles and seconds to hours. From our knowledge of SI prefixes we know that 1 km = 10³ m. From the relationships given on the back inside cover of the book, we find that 1 mi = 1.6093 km.

Given:

Thus, we can convert m to km and then convert km to mi. From our knowledge of time we know that 60 s = 1 min and 60 min = 1 hr. Thus, we can convert s to min and then convert min to hr. The overall process is

Find:

Use Use Use Use m/s

1 km 10³ m

km/s mi/s 1 mi

1.6093 km

60 s

1 min

mi/min mi/hr 60 min

1 hr

Applying first the conversions for distance and then those for time, we can set up one long equation in which unwanted units are canceled:

10³ _{m b a} 1 mi

Speed _{in mi>hr = a515}^m_{s b a} 1 km = 1.15 * 10³ mi>hr

_{1.6093 km b a} 60 s

_{1 minb a}60 min

_{1 hr} b

Our answer has the desired units. We can check our calculation, us ing the estimating procedure described in the “Strategies in Chem istry” box. The given speed is about 500 m>s. Dividing by 1000 converts m to km, giving 0.5 km>s. Because 1 mi is about 1.6 km, this speed corresponds to 0.5>1.6 = 0.3 mi>s. Multiplying by 60 gives about 0.3 * 60 = 20 mi>min. Multiplying again by 60 gives 20 * 60 = 1200 mi>hr. The approximate solution (about 1200 mi/hr) and the detailed solution (1150 mi/hr) are reasonably close. The answer to the detailed solution has three significant figures, cor responding to the number of significant figures in the given speed in m/s.

Conversions Involving Volume

Practice Exercise 1

Fabiola, who lives in Mexico City, fills her car with gas, paying 357 pesos for 40.0 L. What is her fuel cost in dollars per gallon, if 1 peso = 0.0759 dollars? (a) $1.18/gal, (b) $3.03/gal, (c) $1.47/gal, (d) $9.68/gal, (e) $2.56/gal.

Practice Exercise 2

A car travels 28 mi per gallon of gasoline. What is the mileage in kilometers per liter?

The conversion factors previously noted convert from one unit of a given measure to another unit of the same measure, such as from length to length. We also have conver sion factors that convert from one measure to a different one. The density of a sub stance, for example, can be treated as a conversion factor between mass and volume. Suppose we want to know the mass in grams of 2 cubic inches 12.00 in.³2 of gold, which has a density of 19.3 g>cm³. The density gives us the conversion factors:

19.3 g

1 cm³ _and 1 cm³

19.3 g

Because we want a mass in grams, we use the first factor, which has mass in grams in the numerator. To use this factor, however, we must first convert cubic inches to cubic

30 chapter 1 Introduction: Matter and Measurement

centimeters. The relationship between in.³ and cm³ is not given on the back inside

cover, but the relationship between inches and centimeters is given: 1 in. = 2.54 cm

(exactly). Cubing both sides of this equation gives 11 in.2³ = 12.54 cm2³, from which

we write the desired conversion factor:

12.54 cm2³

11 in.2³ ₌ 12.542³ cm³

112³ in.³ ₌ 16.39 cm³

1 in.³

Notice that both the numbers and the units are cubed. Also, because 2.54 is an exact

number, we can retain as many digits of 12.542³ as we need. We have used four, one

more than the number of digits in the density 119.3 g>cm³2. Applying our conversion

factors, we can now solve the problem:

1 in.³ _{b a}19.3 g

Mass in grams = 12.00 in.³_2a16.39 cm³

1 cm³ b = 633 g

The procedure is diagrammed here. The final answer is reported to three significant figures, the same number of significant figures as in 2.00 in.³ and 19.3 g.

Given:

Use

2.54 cm ³

1 in.

in.³

cm³

Find:

Use

19.3 g

1 cm³

g

Sample

Exercise 1.12 Converting Volume Units

Earth’s oceans contain approximately 1.36 * 10⁹ km³ of water. Calculate the volume in liters.

Solution

From the back inside cover, we find 1 L = 10^-3 m³, but there is no relationship listed in volving km³. From our knowledge of SI prefixes, however, we know 1 km = 10³ m and we can use this relationship between lengths to write the desired conversion factor between volumes:

How many liters of water do Earth’s oceans contain?

_a10³ m

3

_{1 km} b

Thus, converting from km³ to m³ to L, we have

₌ 10⁹ m³ 1 km³

Volume in liters = 11.36 * 10⁹ km³_2a10⁹ m³

1 km³ _{b a} 1 L

10^-3 m³ b = 1.36 * 10²¹ L

Practice Exercise 1

A barrel of oil as measured on the oil market is equal to 1.333 U.S. barrels. A U.S. barrel is equal to 31.5 gal. If oil is on the market at $94.0 per barrel, what is the price in dollars per gallon? (a) $2.24/gal, (b) $3.98/gal, (c) $2.98/gal, (d) $1.05/gal, (e) $8.42/gal.

Practice Exercise 2

The surface area of Earth is 510 * 10⁶ km², and 71% of this is ocean. Using the data from the sample exercise, calculate the average depth of the world’s oceans in feet.

Strategies in Chemistry

The Importance of Practice

If you have ever played a musical instrument or participated in ath letics, you know that the keys to success are practice and discipline. You cannot learn to play a piano merely by listening to music, and you cannot learn how to play basketball merely by watching games on television. Likewise, you cannot learn chemistry by merely watch ing your instructor give lectures. Simply reading this book, listening to lectures, or reviewing notes will not usually be sufficient when exam time comes around. Your task is to master chemical concepts and practices to a degree that you can put them to use in solving problems and answering questions. Solving problems correctly takes practice— actually, a fair amount of it. You will do well in your chemistry course if you embrace the idea that you need to master the materials pre sented, and then learn how to apply them in solving problems. Even if you’re a brilliant student, this will take time; it’s what being a stu dent is all about. Almost no one fully absorbs new material on a first reading, especially when new concepts are being presented. You are

Sample

Exercise 1.13 Conversions Involving Density

section 1.6 Dimensional Analysis 31

sure to more fully master the content of the chapters by reading them through at least twice, even more for passages that present you with difficulties in understanding.

Throughout the book, we have provided sample exercises in which the solutions are shown in detail. For practice exercises, we sup ply only the answer, at the back of the book. It is important that you use these exercises to test yourself.

The practice exercises in this text and the homework assignments given by your instructor provide the minimal practice that you will need to succeed in your chemistry course. Only by working all the as signed problems will you face the full range of difficulty and coverage that your instructor expects you to master for exams. There is no sub stitute for a determined and perhaps lengthy effort to work problems on your own. If you are stuck on a problem, however, ask for help from your instructor, a teaching assistant, a tutor, or a fellow student. Spending an inordinate amount of time on a single exercise is rarely effective unless you know that it is particularly challenging and is ex pected to require extensive thought and effort.

What is the mass in grams of 1.00 gal of water? The density of water is 1.00 g/mL. Solution

Before we begin solving this exercise, we note the following:

(1) We are given 1.00 gal of water (the known, or given, quantity) and asked to calculate its mass in grams (the unknown).

(2) We have the following conversion factors either given, commonly known, or available on the back inside cover of the text:

1.00 g water 1 mL water

1 L

1000 mL

1 L

1.057 qt

1 gal 4 qt

The first of these conversion factors must be used as written (with grams in the numerator) to give the desired result, whereas the last conversion factor must be inverted in order to cancel gallons:

Mass _{in grams = 11.00 gal2a} 4 qt _{1 galb a} 1 L

_{1 L b a}1.00 g

_{1 mL} b

= 3.78 * 10³ g water

_{1.057 qtb a}1000 mL

The unit of our final answer is appropriate, and we have taken care of our significant figures. We

can further check our calculation by estimating. We can round 1.057 off to 1. Then focusing on the numbers that do not equal 1 gives 4 * 1000 = 4000 g, in agreement with the detailed calculation.

You should also use common sense to assess the reasonableness of your answer. In this case

we know that most people can lift a gallon of milk with one hand, although it would be tiring to

carry it around all day. Milk is mostly water and will have a density not too different from that of water. Therefore, we might estimate that a gallon of water has mass that is more than 5 lb but less than 50 lb. The mass we have calculated, 3.78 kg * 2.2 lb>kg = 8.3 lb, is thus reasonable as an

order-of-magnitude estimate.

Practice Exercise 1

Trex is a manufactured substitute for wood compounded from post-consumer plastic and wood.

It is frequently used in outdoor decks. Its density is reported as 60 lb>ft³. What is the density of

Trex in kg/L? (a) 138 kg/L, (b) 0.960 kg/L, (c) 259 kg/L, (d) 15.8 kg/L, (e) 11.5 kg/L.

Practice Exercise 2

The density of the organic compound benzene is 0.879 g/mL. Calculate the mass in grams of

1.00 qt of benzene.

A Trex deck.

32 chapter 1 Introduction: Matter and Measurement Strategies in Chemistry

The Features of This Book

If, like most students, you haven’t yet read the part of the Preface to this text entitled TO THE STUDENT, you should do it now. In less than two pages of reading you will encounter valuable advice on how to navigate your way through this book and through the course. We’re serious! This is advice you can use.

The TO THE STUDENT section describes how text features such as “What’s Ahead,” Key Terms, Learning Outcomes, and Key Equations will help you remember what you have learned. We describe there also how to take advantage of the text’s Web site, where many types of online study tools are available. If you have registered for MasteringChemistry®, you will have access to many helpful animations, tutorials, and additional problems correlated to specific topics and sections of each chapter. An in teractive eBook is also available online.

As previously mentioned, working exercises is very important— in fact, essential. You will find a large variety of exercises at the end of each chapter that are designed to test your problem-solving skills in chemistry. Your instructor will very likely assign some of these end-of-chapter exercises as homework. The first few exercises called

Chapter Summary and Key Terms

The Study of Chemistry (Section 1.1) Chemistry is the study of the composition, structure, properties, and changes of matter. The composition of matter relates to the kinds of elements it contains. The structure of matter relates to the ways the atoms of these elements are arranged. A property is any characteristic that gives a sample of mat

ter its unique identity. A molecule is an entity composed of two or more atoms with the atoms attached to one another in a specific way.

Classifications of Matter (Section 1.2) Matter exists in three physical states, gas, liquid, and solid, which are known as the states of matter. There are two kinds of pure substances: elements and compounds. Each element has a single kind of atom and is represented by a chemical symbol consisting of one or two letters, with the first letter capitalized. Compounds are composed of two or more elements joined chemically. The law of constant composition, also called the law of definite proportions, states that the elemental composition of a pure compound is always the same. Most matter consists of a mixture of substances. Mixtures have variable compositions and can be either homogeneous or heterogeneous; homogeneous mixtures are called solutions.

Properties of Matter (Section 1.3) Each substance has a unique set of physical properties and chemical properties that can be used to identify it. During a physical change, matter does not change its com position. Changes of state are physical changes. In a chemical change (chemical reaction) a substance is transformed into a chemically different substance. Intensive properties are independent of the amount of matter examined and are used to identify substances. Extensive properties relate to the amount of substance present. Differences in physical and chemi cal properties are used to separate substances.

The scientific method is a dynamic process used to answer ques tions about the physical world. Observations and experiments lead to tentative explanations or hypotheses. As a hypothesis is tested and re fined, a theory may be developed that can predict the results of future observations and experiments. When observations repeatedly lead to

“Visualizing Concepts” are meant to test how well you understand a concept without plugging a lot of numbers into a formula. The other exercises are grouped in pairs, with the answers given at the back of the book to the odd-numbered exercises (those with red exercise num

bers). An exercise with a [bracket] around its number is designed to be more challenging. Additional Exercises appear after the regular exercises; the chapter sections that they cover are not identified, and they are not paired. Integrative Exercises, which start appearing from Chapter 3, are problems that require skills learned in previous chap

ters. Also first appearing in Chapter 3, are Design an Experiment ex ercises consisting of problem scenarios that challenge you to design experiments to test hypotheses.

Many chemical databases are available, usually on the Web. The CRC Handbook of Chemistry and Physics is the standard reference for many types of data and is available in libraries. The Merck Index is a stan dard reference for the properties of many organic compounds, especially ones of biological interest. WebElements (http://www.webelements .com/) is a good Web site for looking up the properties of the elements. Wolfram Alpha (http://www.wolframalpha.com/) can also be a source of useful information on substances, numerical values, and other data.

the same consistent results, we speak of a scientific law, a general rule that summarizes how nature behaves.

Units of Measurement (Section 1.4) Measurements in chem istry are made using the metric system. Special emphasis is placed on SI units, which are based on the meter, the kilogram, and the second as the basic units of length, mass, and time, respectively. SI units use pre fixes to indicate fractions or multiples of base units. The SI temperature scale is the Kelvin scale, although the Celsius scale is frequently used as well. Absolute zero is the lowest temperature attainable. It has the value 0 K. A derived unit is obtained by multiplication or division of SI base units. Derived units are needed for defined quantities such as speed or volume. Density is an important defined quantity that equals mass divided by volume.

Uncertainty in Measurement (Section 1.5) All measured quantities are inexact to some extent. The precision of a measurement indicates how closely different measurements of a quantity agree with one another. The accuracy of a measurement indicates how well a measurement agrees with the accepted or “true” value. The significant figures in a measured quantity include one estimated digit, the last digit of the measurement. The significant figures indicate the extent of the uncertainty of the measurement. Certain rules must be followed so that a calculation involving measured quantities is reported with the appropriate number of significant figures.

Dimensional Analysis (Section 1.6) In the dimensional analysis approach to problem solving, we keep track of units as we carry measurements through calculations. The units are multiplied together, divided into each other, or canceled like algebraic quantities. Obtaining the proper units for the final result is an important means of checking the method of calculation. When converting units and when carrying out several other types of problems, conversion factors can be used. These factors are ratios constructed from valid relations between equivalent quantities.

Exercises 33 Learning Outcomes After studying this chapter, you should be able to:

• Distinguish among elements, compounds, and mixtures. (Section 1.2) • Identify symbols of common elements. (Section 1.2) • Identify common metric prefixes. (Section 1.4)

Key Equations

• Demonstrate the use of significant figures, scientific notation, and SI units in calculations. (Section 1.5)

• Attach appropriate SI units to defined quantities, and employ dimensional analysis in calculations. (Sections 1.4 and 1.6)

• K = °C + 273.15 [1.1] Converting between Celsius 1°C2 and Kelvin (K) temperature scales _{• °C =} ⁵_{91°F - 322 or °F =} ⁹₅1°C2 + 32 [1.2] Converting between Celsius 1°C2 and Fahrenheit 1°F2 ^tempera ture scales _{• Density =} mass

_volume [1.3] Definition of density

Exercises

Visualizing Concepts

1.1 Which of the following figures represents (a) a pure element, (b) a mixture of two elements, (c) a pure compound, (d) a mixture of an element and a compound? (More than one picture might fit each description.) [Section 1.2]

(i) (ii) (iii)

1.3 Describe the separation method(s) involved in brewing a cup of coffee. [Section 1.3]

1.4 Identify each of the following as measurements of length,

(iv) (v) (vi)

1.2 Does the following diagram represent a chemical or physical change? How do you know? [Section 1.3]

area, volume, mass, density, time, or temperature: (a) 25 ps, (b) 374.2 mg, (c) 77 K, (d) 100,000 km², (e) 1.06 mm, (f) 16 nm², (g) -78 °C, (h) 2.56 g>cm³, (i) 28 cm³. [Section 1.4]

1.5 (a) Three spheres of equal size are composed of aluminum 1density = 2.70 g>cm³2, silver 1density = 10.49 g>cm³2, and nickel 1density = 8.90 g>cm³2. List the spheres from lightest to heaviest. (b) Three cubes of equal mass are composed of gold 1density = 19.32 g>cm³2, platinum 1density = 21.45 g>cm³2, and lead 1density = 11.35 g>cm³2. List the cubes from smallest to largest. [Section 1.4]

1.6 The three targets from a rifle range shown on the next page were produced by: (A) the instructor firing a newly acquired target rifle; (B) the instructor firing his personal target rifle; and (C) a student who has fired his target rifle only a few times. (a) Comment on the accuracy and precision for each of these three sets of results. (b) For the A and C results in the future to look like those in B, what needs to happen? [Section 1.5]

34 chapter 1 Introduction: Matter and Measurement

A B C

1.7 (a) What is the length of the pencil in the following figure if the ruler reads in centimeters? How many significant figures are there in this measurement? (b) An automobile speed ometer with circular scales reading both miles per hour and kilometers per hour is shown. What speed is indicated, in both units? How many significant figures are in the measure ments? [Section 1.5]

1 2 3 4 5 6 7 8 9

1.8 (a) How many significant figures should be reported for the volume of the metal bar shown here? (b) If the mass of the bar is 104.72 g, how many significant figures should be reported when its density is determined using the calculated volume? [Section 1.5]

2.5 cm

1.25 cm

^{5.30 cm}

1.9 When you convert units, how do you decide which part of the conversion factor is in the numerator and which is in the de nominator? [Section 1.6]

1.10 Show the steps to convert the speed of sound, 344 meters per second, into miles per hour. [Section 1.6]

1.11 Consider the jar of jelly beans in the photo. To get an estimate of the number of beans in the jar you weigh six beans and obtain masses of 3.15, 3.12, 2.98, 3.14, 3.02, and 3.09 g. Then you weigh the jar with all the beans in it, and obtain a mass of 2082 g. The empty jar has a mass of 653 g. Based on these data estimate the number of beans in the jar. Justify the number of significant figures you use in your estimate. [Section 1.5]

1.12 The photo below shows a picture of an agate stone. Jack, who picked up the stone on the Lake Superior shoreline and pol ished it, insists that agate is a chemical compound. Ellen ar gues that it cannot be a compound. Discuss the relative merits of their positions. [Section 1.2]

Classification and Properties of Matter (Sections 1.2 and 1.3)

1.13 Classify each of the following as a pure substance or a mixture. If a mixture, indicate whether it is homogeneous or hetero geneous: (a) rice pudding, (b) seawater, (c) magnesium, (d) crushed ice.

1.14 Classify each of the following as a pure substance or a mixture. If a mixture, indicate whether it is homogeneous or heterogeneous: (a) air, (b) tomato juice, (c) iodine crystals, (d) sand.

1.15 Give the chemical symbol or name for the following elements, as appropriate: (a) sulfur, (b) gold, (c) potassium, (d) chlorine, (e) copper, (f) U, (g) Ni, (h) Na, (i) Al, (j) Si.

1.16 Give the chemical symbol or name for each of the follow ing elements, as appropriate: (a) carbon, (b) nitrogen, (c) titanium, (d) zinc, (e) iron, (f) P, (g) Ca, (h) He, (i) Pb, (j) Ag.

1.17 A solid white substance A is heated strongly in the absence of air. It decomposes to form a new white substance B and a gas C. The gas has exactly the same properties as the prod uct obtained when carbon is burned in an excess of oxygen. Based on these observations, can we determine whether solids A and B and gas C are elements or compounds? Explain your conclusions for each substance.

1.18 You are hiking in the mountains and find a shiny gold nug get. It might be the element gold, or it might be “fool’s gold,” which is a nickname for iron pyrite, FeS₂. What kinds of ex periments could be done to determine if the shiny nugget is really gold?

Exercises 35

1.19 In the process of attempting to characterize a substance, a chemist makes the following observations: The sub stance is a silvery white, lustrous metal. It melts at 649 °C and boils at 1105 °C. Its density at 20 °C is 1.738 g>cm³. The substance burns in air, producing an intense white light. It reacts with chlorine to give a brittle white solid. The substance can be pounded into thin sheets or drawn into wires. It is a good conductor of electricity. Which of these characteristics are physical properties, and which are chemical properties?

1.20 (a) Read the following description of the element zinc and in dicate which are physical properties and which are chemical properties.

Zinc melts at 420 °C. When zinc granules are added to dilute sulfuric acid, hydrogen is given off and the metal dissolves. Zinc has a hardness on the Mohs scale of 2.5 and a density of 7.13g>cm³ at 25 °C. It reacts slowly with oxygen gas at el

evated temperatures to form zinc oxide, ZnO.

(b) Which properties of zinc can you describe from the photo? Are these physical or chemical properties?

1.21 Label each of the following as either a physical process or a chemical process: (a) rusting of a metal can, (b) boiling a cup of water, (c) pulverizing an aspirin, (d) digesting a candy bar, (e) exploding of nitroglyerin.

1.22 A match is lit and held under a cold piece of metal. The following observations are made: (a) The match burns. (b) The metal gets warmer. (c) Water condenses on the metal. (d) Soot (carbon) is deposited on the metal. Which of these occurrences are due to physical changes, and which are due to chemical changes?

1.23 Suggest a method of separating each of the following mixtures into two components: (a) sugar and sand, (b) oil and vinegar. 1.24 Three beakers contain clear, colorless liquids. One beaker contains pure water, another contains salt water, and an other contains sugar water. How can you tell which beaker is which? (No tasting allowed!)

Units and Measurement (Section 1.4)

1.25 What exponential notation do the following abbreviations represent? (a) d, (b) c, (c) f, (d) m, (e) M, (f) k, (g) n, (h) m, (i) p.

1.26 Use appropriate metric prefixes to write the following mea surements without use of exponents: (a) 2.3 * 10^-10 L, (b) 4.7 * 10^-6 g, (c) 1.85 * 10^-12 m, (d) 16.7 * 10⁶ s, (e) 15.7 * 10³ g, (f) 1.34 * 10^-3 m, (g) 1.84 * 10² cm.

1.27 Make the following conversions: (a) 72 °F to °C, (b) 216.7 °C to °F, (c) 233 °C to K, (d) 315 K to °F, (e) 2500 °F to K, (f) 0 K to °F.

1.28 (a) The temperature on a warm summer day is 87 °F. What is the temperature in °C? (b) Many scientific data are reported at 25 °C. What is this temperature in kelvins and in degrees Fahrenheit? (c) Suppose that a recipe calls for an oven temperature of 400 °F. Convert this temperature to degrees Celsius and to kelvins. (d) Liquid nitrogen boils at 77 K. Convert this temperature to degrees Fahrenheit and to degrees Celsius.

1.29 (a) A sample of tetrachloroethylene, a liquid used in dry cleaning that is being phased out because of its potential to cause cancer, has a mass of 40.55 g and a volume of 25.0 mL at 25 °C. What is its density at this temperature? Will tetra

chloroethylene float on water? (Materials that are less dense than water will float.) (b) Carbon dioxide 1CO₂2 is a gas at room temperature and pressure. However, carbon dioxide can be put under pressure to become a “supercritical fluid” that is a much safer dry-cleaning agent than tetrachloroethyl

ene. At a certain pressure, the density of supercritical CO₂ is 0.469 g>cm³. What is the mass of a 25.0-mL sample of super critical CO₂ at this pressure?

1.30 (a) A cube of osmium metal 1.500 cm on a side has a mass of 76.31 g at 25 °C. What is its density in g>cm³ at this tempera ture? (b) The density of titanium metal is 4.51g>cm³ at 25 °C. What mass of titanium displaces 125.0 mL of water at 25 °C? (c) The density of benzene at 15 °C is 0.8787g>mL. Calculate the mass of 0.1500 L of benzene at this temperature.

1.31 (a) To identify a liquid substance, a student determined its density. Using a graduated cylinder, she measured out a 45-mL sample of the substance. She then measured the mass of the sample, finding that it weighed 38.5 g. She knew that the substance had to be either isopropyl alcohol 1density 0.785 g>mL2 or toluene 1density 0.866>mL2. What are the calculated density and the probable identity of the substance? (b) An experiment requires 45.0 g of ethylene gly col, a liquid whose density is 1.114 g>mL. Rather than weigh the sample on a balance, a chemist chooses to dispense the liquid using a graduated cylinder. What volume of the liquid should he use? (c) Is a graduated cylinder such as that shown in Figure 1.19 likely to afford the accuracy of measurement needed? (d) A cubic piece of metal measures 5.00 cm on each edge. If the metal is nickel, whose density is 8.90 g>cm³, what is the mass of the cube?

1.32 (a) After the label fell off a bottle containing a clear liquid be lieved to be benzene, a chemist measured the density of the liquid to verify its identity. A 25.0-mL portion of the liquid had a mass of 21.95 g. A chemistry handbook lists the den sity of benzene at 15 °C as 0.8787 g>mL. Is the calculated density in agreement with the tabulated value? (b) An experi ment requires 15.0 g of cyclohexane, whose density at 25 °C is 0.7781 g>mL. What volume of cyclohexane should be used? (c) A spherical ball of lead has a diameter of 5.0 cm. What is the mass of the sphere if lead has a density of 11.34 g>cm³? (The volume of a sphere is 14>32pr³, where r is the radius.)

1.33 In the year 2011, an estimated amount of 35 billion tons of carbon dioxide 1CO₂2 was emitted worldwide due to fossil fuel combustion and cement production. Express this mass of CO₂ in grams without exponential notation, using an appro

priate metric prefix.

36 chapter 1 Introduction: Matter and Measurement

1.34 Silicon for computer chips is grown in large cylinders called “boules” that are 300 mm in diameter and 2 m in length, as shown. The density of silicon is 2.33 g>cm³. Silicon wafers for making integrated circuits are sliced from a 2.0 m boule and are typically 0.75 mm thick and 300 mm in diameter. (a) How many wafers can be cut from a single boule? (b) What is the mass of a silicon wafer? (The volume of a cylinder is given by pr²h, where r is the radius and h is its height.)

(a) 320.5 - 16104.5>2.32

(b) 31285.3 * 10⁵2 - 11.200 * 10³24 * 2.8954

(c) 10.0045 * 20,000.02 + 12813 * 122

(d) 863 * 31255 - 13.45 * 10824

1.43 You weigh an object on a balance and read the mass in grams according to the picture. How many significant figures are in this measurement?

Diamond blade Si boule

2 m

0.75 mm

thickness

300 mm

diameter

Cut wafers

1.44 You have a graduated cylinder that contains a liquid (see pho

Uncertainty in Measurement (Section 1.5)

1.35 Indicate which of the following are exact numbers: (a) the mass of a 3 by 5–inch index card, (b) the number of ounces in a pound, (c) the volume of a cup of Seattle’s Best coffee, (d) the number of inches in a mile, (e) the number of micro

seconds in a week, (f) the number of pages in this book. 1.36 Indicate which of the following are exact numbers: (a) the mass of a 32-oz can of coffee, (b) the number of students in your chemistry class, (c) the temperature of the surface of the Sun, (d) the mass of a postage stamp, (e) the number of mil liliters in a cubic meter of water, (f) the average height of NBA basketball players.

1.37 What is the number of significant figures in each of the fol lowing measured quantities? (a) 601 kg, (b) 0.054 s, (c) 6.3050 cm, (d) 0.0105 L, (e) 7.0500 * 10^-3 m³, (f) 400 g.

1.38 Indicate the number of significant figures in each of the following measured quantities: (a) 3.774 km, (b) 205 m², (c) 1.700 cm, (d) 350.00 K, (e) 307.080 g, (f) 1.3 * 10³ m>s.

1.39 Round each of the following numbers to four significant fig ures and express the result in standard exponential notation: (a) 102.53070, (b) 656.980, (c) 0.008543210, (d) 0.000257870, (e) -0.0357202.

1.40 (a) The diameter of Earth at the equator is 7926.381 mi. Round this number to three significant figures and express it in stan dard exponential notation. (b) The circumference of Earth through the poles is 40,008 km. Round this number to four sig nificant figures and express it in standard exponential notation.

1.41 Carry out the following operations and express the answers with the appropriate number of significant figures.

(a) 14.3505 + 2.65

(b) 952.7 - 140.7389

(c) 13.29 * 10⁴210.25012

(d) 0.0588/0.677

1.42 Carry out the following operations and express the answer with the appropriate number of significant figures.

tograph). Write the volume of the liquid, in milliliters, using the proper number of significant figures.

Dimensional Analysis (Section 1.6)

1.45 Using your knowledge of metric units, English units, and the information on the back inside cover, write down the conver sion factors needed to convert (a) mm to nm, (b) mg to kg, (c) km to ft, (d) in.³ to cm³.

1.46 Using your knowledge of metric units, English units, and the information on the back inside cover, write down the conver sion factors needed to convert (a) mm to mm, (b) ms to ns, (c) mi to km, (d) ft³ to L.

1.47 (a) A bumblebee flies with a ground speed of 15.2 m/s. Cal culate its speed in km/hr. (b) The lung capacity of the blue whale is 5.0 * 10³ L. Convert this volume into gallons. (c) The Statue of Liberty is 151 ft tall. Calculate its height in meters. (d) Bamboo can grow up to 60.0 cm/day. Convert this growth rate into inches per hour.

1.48 (a) The speed of light in a vacuum is 2.998 * 10⁸ m>s. Calculate its speed in miles per hour. (b) The Sears Tower in Chicago is 1454 ft tall. Calculate its height in meters. (c) The Vehicle Assembly Building at the Kennedy Space Center in Florida has a volume of 3,666,500 m³. Convert this volume to liters and express the result in standard exponential notation. (d) An individual suffering from a high cholesterol level in her

blood has 242 mg of cholesterol per 100 mL of blood. If the total blood volume of the individual is 5.2 L, how many grams of total blood cholesterol does the individual’s body contain?

1.49 The inside dimension of a box that is cubic is 24.8 cm on each edge with an uncertainty of 0.2 cm. What is the volume of the box? What do you estimate to be the uncertainty in the calcu lated volume?

1.50 The distance from Grand Rapids, Michigan, to Detroit is listed in a road atlas as 153 miles. Describe some of the factors that contribute to the uncertainty in this number. To make the num ber more precise, what would you need to specify and measure?

1.51 Perform the following conversions: (a) 5.00 days to s, (b) 0.0550 mi to m, (c) $1.89/gal to dollars per liter, (d) 0.510 in./ms to km/hr, (e) 22.50 gal/min to L/s, (f) 0.02500 ft³ to cm³.

1.52 Carry out the following conversions: (a) 0.105 in. to mm, (b) 0.650 qt to mL, (c) 8.75 mm>s to km>hr,(d) 1.955 m³ to yd³, (e) $3.99/lb to dollars per kg, (f) 8.75 lb>ft³ to g>mL.

1.53 (a) How many liters of wine can be held in a wine barrel whose capacity is 31 gal? (b) The recommended adult dose of Elixophyllin®, a drug used to treat asthma, is 6 mg/kg of body mass. Calculate the dose in milligrams for a 185-lb person. (c) If an automobile is able to travel 400 km on 47.3 L of gaso

line, what is the gas mileage in miles per gallon? (d) When the coffee is brewed according to directions, a pound of coffee beans yields 50 cups of coffee 14 cups = 1 qt2. How many kg of coffee are required to produce 200 cups of coffee?

Additional Exercises

1.59 (a) Classify each of the following as a pure substance, a solu tion, or a heterogeneous mixture: a gold coin, a cup of coffee, a wood plank. (b) What ambiguities are there in answering part (a) from the descriptions given?

1.60 (a) What is the difference between a hypothesis and a theory? (b) Explain the difference between a theory and a scientific law. Which addresses how matter behaves, and which ad dresses why it behaves that way?

1.61 A sample of ascorbic acid (vitamin C) is synthesized in the laboratory. It contains 1.50 g of carbon and 2.00 g of oxy gen. Another sample of ascorbic acid isolated from citrus fruits contains 6.35 g of carbon. How many grams of oxygen does it contain? Which law are you assuming in answering this question?

1.62 Ethyl chloride is sold as a liquid (see photo) under pres sure for use as a local skin anesthetic. Ethyl chloride boils at 12 °C at atmospheric pressure. When the liquid is sprayed onto the skin, it boils off, cooling and numbing the skin as it

vaporizes. (a) What changes of state are involved in this use of ethyl chloride? (b) What is the boiling point of ethyl chlo ride in degrees Fahrenheit? (c) The bottle shown contains 103.5 mL of ethyl chloride. The density of ethyl chloride at 25 °C is 0.765 g>cm³. What is the mass of ethyl chloride in the bottle?

Additional Exercises 37

1.54 (a) If an electric car is capable of going 225 km on a single charge, how many charges will it need to travel from Seattle, Washing ton, to San Diego, California, a distance of 1257 mi, assuming that the trip begins with a full charge? (b) If a migrating loon flies at an average speed of 14 m/s, what is its average speed in mi/hr? (c) What is the engine piston displacement in liters of an engine whose displacement is listed as 450 in.³? (d) In March 1989 the Exxon Valdez ran aground and spilled 240,000 barrels of crude petroleum off the coast of Alaska. One barrel of petroleum is equal to 42 gal. How many liters of petroleum were spilled?

1.55 The density of air at ordinary atmospheric pressure and 25 °C is 1.19 g>L. What is the mass, in kilograms, of the air in a room that measures 14.5 ft * 16.5 ft * 8.0 ft?

1.56 The concentration of carbon monoxide in an urban apart ment is 48 mg>m³. What mass of carbon monoxide in grams is present in a room measuring 10.6 ft * 14.8 ft * 20.5 ft?

1.57 Gold can be hammered into extremely thin sheets called gold leaf. An architect wants to cover a 100 ft * 82 ft ceiling with gold leaf that is five–millionths of an inch thick. The density of gold is 19.32 g>cm³, and gold costs $1654 per troy ounce 11 troy ounce = 31.1034768 g2. How much will it cost the architect to buy the necessary gold?

1.58 A copper refinery produces a copper ingot weighing 150 lb. If the copper is drawn into wire whose diameter is 7.50 mm, how many feet of copper can be obtained from the ingot? The density of copper is 8.94 g>cm³. (Assume that the wire is a cylinder whose volume V = pr²h, where r is its radius and h

is its height or length.)

1.63 Two students determine the percentage of lead in a sample as a laboratory exercise. The true percentage is 22.52%. The

students’ results for three determinations are as follows: (1) 22.52, 22.48, 22.54

(2) 22.64, 22.58, 22.62

38 chapter 1 Introduction: Matter and Measurement

(a) Calculate the average percentage for each set of data and state which set is the more accurate based on the average. (b) Precision can be judged by examining the average of the deviations from the average value for that data set. (Calculate the average value for each data set; then calculate the average value of the absolute deviations of each measurement from the average.) Which set is more precise?

1.64 Is the use of significant figures in each of the following statements appropriate? Why or why not? (a) Apple sold 22,727,000 iPods during the last three months of 2008. (b) New York City receives 49.7 inches of rain, on average, per year. (c) In the United States, 0.621% of the population has the surname Brown. (d) You calculate your grade point average to be 3.87562.

1.65 What type of quantity (for example, length, volume, density) do the following units indicate? (a) mL, (b) cm², (c) mm³, (d) mg/L, (e) ps, (f) nm, (g) K.

1.66 Give the derived SI units for each of the following quantities in base SI units:

(a) acceleration = distance>time²

(b) force = mass * acceleration

(c) work = force * distance

(d) pressure = force>area

(e) power = work>time

(f) velocity = distance>time

(g) energy = mass * 1velocity2²

1.67 The distance from Earth to the Moon is approximately 240,000 mi. (a) What is this distance in meters? (b) The per egrine falcon has been measured as traveling up to 350 km/ hr in a dive. If this falcon could fly to the Moon at this speed, how many seconds would it take? (c) The speed of light is 3.00 * 10⁸ m>s. How long does it take for light to travel from Earth to the Moon and back again? (d) Earth travels around the Sun at an average speed of 29.783 km>s. Convert this speed to miles per hour.

1.68 Which of the following would you characterize as a pure or nearly pure substance? (a) baking powder; (b) lemon juice; (c) propane gas, used in outdoor gas grills; (d) aluminum foil; (e) ibuprofen; (f) bourbon whiskey; (g) helium gas; (h) clear water pumped from a deep aquifer.

1.69 The U.S. quarter has a mass of 5.67 g and is approximately 1.55 mm thick. (a) How many quarters would have to be stacked to reach 575 ft, the height of the Washing ton Monument? (b) How much would this stack weigh? (c) How much money would this stack contain? (d) The U.S. National Debt Clock showed the outstanding public debt to be $16,213,166,914,811 on October 28, 2012. How many stacks like the one described would be necessary to pay off this debt?

1.70 In the United States, water used for irrigation is measured in acre-feet. An acre-foot of water covers an acre to a depth of exactly 1 ft. An acre is 4840 yd². An acre-foot is enough water to supply two typical households for 1.00 yr. (a) If desalinated water costs $1950 per acre-foot, how much does desalinated water cost per liter? (b) How much would it cost one house

hold per day if it were the only source of water?

1.71 By using estimation techniques, determine which of the follow ing is the heaviest and which is the lightest: a 5-lb bag of potatoes, a 5-kg bag of sugar, or 1 gal of water 1density = 1.0 g>mL2.

1.72 Suppose you decide to define your own temperature scale with units of O, using the freezing point 113 °C2 and boiling point 1360 °C2 of oleic acid, the main component of olive oil. If you set the freezing point of oleic acid as 0 °O and the boiling point as 100 °O, what is the freezing point of water on this new scale?

1.73 The liquid substances mercury 1density = 13.6 g>mL2, water 11.00 g>mL2, and cyclohexane 10.778 g>mL2 do not form a solution when mixed but separate in distinct layers. Sketch how the liquids would position themselves in a test tube.

1.74 Two spheres of equal volume are placed on the scales as shown. Which one is more dense?

1.75 Water has a density of 0.997 g>cm³ at 25 °C; ice has a density of 0.917 g>cm³ at -10 °C. (a) If a soft-drink bottle whose vol ume is 1.50 L is completely filled with water and then frozen to -10 °C, what volume does the ice occupy? (b) Can the ice be contained within the bottle?

1.76 A 32.65-g sample of a solid is placed in a flask. Toluene, in which the solid is insoluble, is added to the flask so that the total volume of solid and liquid together is 50.00 mL. The solid and toluene together weigh 58.58 g. The density of toluene at the temperature of the experiment is 0.864 g>mL.

What is the density of the solid?

1.77 A thief plans to steal a gold sphere with a radius of 28.9 cm from a museum. If the gold has a density of 19.3 g>cm³, what is the mass of the sphere in pounds? [The volume of a sphere is V = 14>32pr³.4 Is the thief likely to be able to walk off with the gold sphere unassisted?

1.78 Automobile batteries contain sulfuric acid, which is com monly referred to as “battery acid.” Calculate the number of grams of sulfuric acid in 1.00 gal of battery acid if the solution has a density of 1.28 g/mL and is 38.1% sulfuric acid by mass.

1.79 A 40-lb container of peat moss measures 14 * 20 * 30 in. A 40-lb container of topsoil has a volume of 1.9 gal. (a) Calculate the average densities of peat moss and topsoil in units of g>cm³. Would it be correct to say that peat moss is “lighter” than topsoil? Explain. (b) How many bags of peat moss are needed to cover an area measuring 15.0 ft * 20.0 ft to a depth of 3.0 in.?

1.80 A package of aluminum foil contains 50 ft² of foil, which weighs approximately 8.0 oz. Aluminum has a density of 2.70 g>cm³. What is the approximate thickness of the foil in millimeters?

1.81 The total rate at which power used by humans worldwide is approximately 15 TW (terawatts). The solar flux aver aged over the sunlit half of Earth is 680 W>m². (assuming no clouds). The area of Earth’s disc as seen from the sun is 1.28 * 10¹⁴ m². The surface area of Earth is approxi mately 197,000,000 square miles. How much of Earth’s

surface would we need to cover with solar energy collectors to power the planet for use by all humans? Assume that the solar energy collectors can convert only 10% of the available sunlight into useful power.

1.82 In 2005, J. Robin Warren and Barry J. Marshall shared the Nobel Prize in Medicine for discovery of the bacterium Helicobacter pylori, and for establishing experimental proof that it plays a major role in gastritis and peptic ulcer disease. The story began when Warren, a pathologist, noticed that bacilli were associated with the tissues taken from patients suffering from ulcers. Look up the history of this case and describe Warren’s first hypothesis. What sorts of evidence did it take to create a credible theory based on it?

1.83 A 25.0-cm long cylindrical glass tube, sealed at one end, is filled with ethanol. The mass of ethanol needed to fill the tube is found to be 45.23 g. The density of ethanol is 0.789 g/mL. Calculate the inner diameter of the tube in centimeters.

1.84 Gold is alloyed (mixed) with other metals to increase its hard ness in making jewelry. (a) Consider a piece of gold jewelry that weighs 9.85 g and has a volume of 0.675 cm³. The jew elry contains only gold and silver, which have densities of 19.3 and 10.5 g>cm³, respectively. If the total volume of the jewelry is the sum of the volumes of the gold and silver that it contains, calculate the percentage of gold (by mass) in the jewelry. (b) The relative amount of gold in an alloy is com monly expressed in units of carats. Pure gold is 24 carat, and the percentage of gold in an alloy is given as a percentage of this value. For example, an alloy that is 50% gold is 12 carat. State the purity of the gold jewelry in carats.

1.85 Paper chromatography is a simple but reliable method for sep arating a mixture into its constituent substances. You have a mixture of two vegetable dyes, one red and one blue, that you are trying to separate. You try two different chromatography procedures and achieve the separations shown in the figure. Which procedure worked better? Can you suggest a method to quantify how good or poor the separation was?

Additional Exercises 39

1.86 Judge the following statements as true or false. If you believe a statement to be false, provide a corrected version.

(a) Air and water are both elements.

(b) All mixtures contain at least one element and one compound.

(c) Compounds can be decomposed into two or more other substances; elements cannot.

(d) Elements can exist in any of the three states of matter. (e) When yellow stains in a kitchen sink are treated with bleach water, the disappearance of the stains is due to a physical change.

(f) A hypothesis is more weakly supported by experimental evidence than a theory.

(g) The number 0.0033 has more significant figures than 0.033.

(h) Conversion factors used in converting units always have a numerical value of one.

(i) Compounds always contain at least two different elements.

1.87 You are assigned the task of separating a desired granular ma terial with a density of 3.62 g>cm³ from an undesired granular material that has a density of 2.04 g>cm³. You want to do this by shaking the mixture in a liquid in which the heavier mate rial will fall to the bottom and the lighter material will float. A solid will float on any liquid that is more dense. Using an Internet-based source or a handbook of chemistry, find the densities of the following substances: carbon tetrachloride, hexane, benzene, and diiodomethane. Which of these liquids will serve your purpose, assuming no chemical interaction be tween the liquid and the solids?

1.88 In 2009, a team from Northwestern University and Western Washington University reported the preparation of a new “spongy” material composed of nickel, molybdenum, and sulfur that excels at removing mercury from water. The den

sity of this new material is 0.20 g>cm³, and its surface area is 1242 m² per gram of material. (a) Calculate the volume of a 10.0-mg sample of this material. (b) Calculate the surface area for a 10.0-mg sample of this material. (c) A 10.0-mL sample of contaminated water had 7.748 mg of mercury in it. After treatment with 10.0 mg of the new spongy material, 0.001 mg of mercury remained in the contaminated water. What percentage of the mercury was removed from the water? (d) What is the final mass of the spongy material after the exposure to mercury?

2

Atoms, Molecules, and Ions

Look around at the great variety of colors, textures, and other properties in the materials that surround you—the colors in a garden, the texture of the fabric in your clothes, the solubility of sugar in a cup of coffee, or the beauty and complexity of a geode like the one shown to the right. How can we explain the striking and seemingly infinite variety of properties of the materials that make up our world? What makes diamonds transparent and hard? A large crystal of sodium chloride, table salt, looks a bit like a diamond, but is brittle and readily dissolves in water. What accounts for the differences? Why does paper burn, and why does water quench fires? The answers to all such questions lie in the structures of atoms, which determine the physical and chemical properties of matter.

Although the materials in our world vary greatly in their properties, everything is formed from only about 100 elements and, therefore, from only about 100 chemically

different kinds of atoms. In a sense, these different atoms are like the 26 letters of the English alphabet that join in different combinations to form the immense number of words in our language. But what rules govern the ways in which atoms combine? How do the properties of a substance relate to the kinds of atoms it contains? Indeed, what is an atom like, and what makes the atoms of one element different from those of another?

In this chapter we introduce the basic structure of atoms and discuss the forma tion of molecules and ions, thereby providing a foundation for exploring chemistry more deeply in later chapters.

What’s

▶ A section through a geode. A geode is a mass of mineral matter (often containing quartz) that accumulates slowly within the shell of a roughly spherical, hollow rock. Eventually, perfectly formed crystals may develop at a geode’s center. The colors of a geode depend upon its composition. Here, agate crystallized out as the geode formed.

Ahead

2.1 The Atomic Theory of Matter We begin with a brief history of the notion of atoms—the smallest pieces of matter.

2.2 The Discovery of Atomic Structure We then look at some key experiments that led to the discovery of electrons and to the nuclear model of the atom.

2.3 The Modern View of Atomic Structure We explore the modern theory of atomic structure, including the ideas of atomic numbers, mass numbers, and isotopes.

2.4 Atomic Weights We introduce the concept of atomic weights and how they relate to the masses of individual atoms.

2.5 The Periodic Table We examine the organization of the periodic table, in which elements are put in order of increasing atomic number and grouped by chemical similarity.

2.6 Molecules and Molecular Compounds We discuss the assemblies of atoms called molecules and how their compositions are represented by empirical and molecular formulas.

2.7 Ions and Ionic Compounds We learn that atoms can gain or lose electrons to form ions. We also look at how to use the periodic table to predict the charges on ions and the empirical formulas of ionic compounds.

2.8 Naming Inorganic Compounds We consider the systematic way in which substances are named, called nomenclature, and how this nomenclature is applied to inorganic compounds.

2.9 Some Simple Organic Compounds We introduce organic chemistry, the chemistry of the element carbon.