Brown LeMay 13e - Chemistry - The Central Science Textbook - AP Chem
Chemistry The Cen t ral S c ien c e 13TH Edition
Chemistry The Cen t ral S c ien c e 13TH 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 Kc 634 Equilibrium Constants in Terms of Pressure, Kp 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 Ka from pH 688 Percent
Ionization 689 Using Ka to Calculate pH 690 Polyprotic Acids 694
16.7 Weak Bases 696
Types of Weak Bases 698
16.8 Relationship between Ka and Kb 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, Ksp 748
Solubility and Ksp 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 Authors 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
The 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
Statea Liquid Gas Gas
Normal boiling point 100 °C -253 °C -183 °C
Densitya 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 solvent
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 1015 1 petawatt (PW) = 1 * 1015 wattsa
Tera T 1012 1 terawatt (TW) = 1 * 1012 watts
Giga G 109 1 gigawatt (GW) = 1 * 109 watts
Mega M 106 1 megawatt (MW) = 1 * 106 watts
Kilo k 103 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 mb 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-intervals
100 °C
100 degree-intervals
212 °F 98.6 °F
Water boils
180 degree-intervals
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 = 59 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 lb2a453.6 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 = a515ms b a 1 km = 1.15 * 103 mi>hr
1.6093 km b a 60 s
1 minb a60 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 a19.3 g
Mass in grams = 12.00 in.32a16.39 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 km32a109 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 a1.00 g
1 mL b
= 3.78 * 103 g water
1.057 qtb a1000 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 = 591°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.
Chemistry The Cen t ral S c ien c e 13TH Edition
Chemistry The Cen t ral S c ien c e 13TH 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 Kc 634 Equilibrium Constants in Terms of Pressure, Kp 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 Ka from pH 688 Percent
Ionization 689 Using Ka to Calculate pH 690 Polyprotic Acids 694
16.7 Weak Bases 696
Types of Weak Bases 698
16.8 Relationship between Ka and Kb 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, Ksp 748
Solubility and Ksp 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
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Preface xxv
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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 Authors 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
The 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
Statea Liquid Gas Gas
Normal boiling point 100 °C -253 °C -183 °C
Densitya 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 solvent
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 1015 1 petawatt (PW) = 1 * 1015 wattsa
Tera T 1012 1 terawatt (TW) = 1 * 1012 watts
Giga G 109 1 gigawatt (GW) = 1 * 109 watts
Mega M 106 1 megawatt (MW) = 1 * 106 watts
Kilo k 103 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 mb 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-intervals
100 °C
100 degree-intervals
212 °F 98.6 °F
Water boils
180 degree-intervals
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 = 59 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 lb2a453.6 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 = a515ms b a 1 km = 1.15 * 103 mi>hr
1.6093 km b a 60 s
1 minb a60 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 a19.3 g
Mass in grams = 12.00 in.32a16.39 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 km32a109 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 a1.00 g
1 mL b
= 3.78 * 103 g water
1.057 qtb a1000 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 = 591°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.