Main Basic Concepts of Chemistry, Eighth Edition

Basic Concepts of Chemistry, Eighth Edition

,
Engineers who need to have a better understanding of chemistry will benefit from this accessible book. It places a stronger emphasis on outcomes assessment, which is the driving force for many of the new features. Each section focuses on the development and assessment of one or two specific objectives. Within each section, a specific objective is included, an anticipatory set to orient the reader, content discussion from established authors, and guided practice problems for relevant objectives. These features are followed by a set of independent practice problems. The expanded Making it Real feature showcases topics of current interest relating to the subject at hand such as chemical forensics and more medical related topics. Numerous worked examples in the text now include Analysis and Synthesis sections, which allow engineers to explore concepts in greater depth, and discuss outside relevance.
Year: 2008
Edition: 8th
Publisher: Wiley
Language: english
Pages: 726
ISBN 13: 9780471741541
ISBN: 047174154X
File: PDF, 22.96 MB
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Sepultura - Just the Riffs

発行年: 1995
言語: english
File: PDF, 6.46 MB
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BASIC CONCE PTS
OF CH E M ISTRY

8th
EDITION

L E O J . M A LO N E

St. Louis University

T H E O D O R E O . D O LT E R Southwestern Illinois College

John Wiley & Sons, Inc.

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This volume contains selected illustrations from the following
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Changes, Fourth Edition, ©2004.
• Brady, James E.; Senese, Fred, Chemistry: Matter and Its
Changes, Fifth Edition, ©2009.
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• Olmsted III, John; Williams, Gregory M., Chemistry, Fourth
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Library of Congress Cataloging-in-Publication Data
Malone Leo J., 1938Basic concepts of chemistry / Leo J. Malone, Theodore O. Dolter. – 8th ed.
p. cm.
ISBN 978-0-471-74154-1 (cloth)
1. Chemistry–Textbooks. I. Dolter, Theodore O., 1965- II. Title.
QD31.3.M344 2010
540–dc22
2008036140
Printed in the United States of America
10 9 8 7 6 5 4 3 2 1

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A B O U T

T H E

A U T H O R S

LEO J. MALON E

Leo Malone is a native of Kansas where he received his B.S. in Chemistry from
Wichita State University in 1960 and M.S. in Chemistry in 1962. At WSU he worked
under the direction of Dr. Robert Christian. He moved on to the University of
Michigan where he received his Ph.D. in 1964 under the direction of Dr. Robert
Parry. Dr. Malone began his teaching career at Saint Louis University in 1965 where
he remained until his retirement as Professor Emeritus in 2005. Although his early
research at SLU involved boron hydride chemistry, he eventually concentrated his
efforts on the teaching of basic chemistry and in the field of chemical education.
T H E O D O R E ( T E D ) O . D O LT E R

Ted Dolter received his B.S. in Chemistry from St. Louis University in 1987, where
he was a student of Dr. Malone’s. He went on to the University of Illinois where he
received a Masters of Chemical Education in 1990. He concurrently earned a
secondary teaching certificate, and it was there that he received most of his training
in modern educational theory. After six years of teaching high school chemistry
and evening courses at Southwestern Illinois College, he joined the faculty at SWIC
full time. In 2004, he was elected to the Department Chair position, where he
currently serves and is involved in the department’s outcomes assessment initiative.
Professor Dolter’s background in educational training and his knowledge of
the needs of the growing numbers of community college students compliment
Dr. Malone’s years of traditional university educational experience. Together, they
have produced a text that remains flexible and applicable to the rapidly changing
face of today’s post-secondary student population.

vii

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B R I E F

PROLOGUE
CHAPTER 1
CHAPTER 2
CHAPTER 3
CHAPTER 4
CHAPTER 5
CHAPTER 6
CHAPTER 7
CHAPTER 8
CHAPTER 9

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C O N T E N T S

Introduction to the Study of Chemistry 2
Measurements in Chemistry
14
Elements and Compounds 50
The Properties of Matter and Energy 78
The Periodic Table and Chemical Nomenclature 108
Chemical Reactions 136
Quantities in Chemistry 168
Quantitative Relationships in Chemical Reactions 196
Modern Atomic Theory
220
The Chemical Bond 252

The Gaseous State
292
CHAPTER 11
The Solid and Liquid States 330
CHAPTER 12
Aqueous Solutions
362
CHAPTER 13
Acids, Bases, and Salts
396
CHAPTER 14
Oxidation–Reduction Reactions 436
CHAPTER 15
Reaction Rates and Equilibrium 470
CHAPTER 16
Nuclear Chemistry 510
CHAPTER 17
Organic Chemistry (WEB)
w-2
CHAPTER 18
Biochemistry (WEB)
w-42
Foreword to the Appendices A-1
APPENDIX A
Basic Mathematics
A-1
APPENDIX B
Basic Algebra
A-10
APPENDIX C
Scientific Notation A-19
APPENDIX D
Graphs A-27
APPENDIX E
Calculators
A-31
APPENDIX F
Glossary
A-36
APPENDIX G
Answers to Problems
A-45
CHAPTER 10

viii

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C O N T E N T S

PROLOGUE

Science and the Magnificent Human Mind
A
B
C

2

The Origin of Matter 4
The Mystery of Fire 9
The Scientific Method 10

CHAPTER 1

Measurements in Chemistry
PA R T

A

14

1-1 The Numerical Value of a Measurement 16
1-2 Significant Figures and Mathematical Operations 19
1-3 Expressing Large and Small Numbers: Scientific Notation
MAKING IT REAL:

Ted Williams and Significant Figures

PA R T A S U M M A R Y

PA R T

B

16

THE NUMBERS USED IN CHEMISTRY

22
22

25

THE MEASUREMENTS USED
IN CHEMISTRY
26

1-4 Measurement of Mass, Length, and Volume 26
1-5 Conversion of Units by the Factor-Label Method 30
MAKING IT REAL:

Worlds from the Small to the Distant—
Picometers to Terameters
31

1-6 Measurement of Temperature 38
PA R T B S U M M A R Y
41
Chapter Summary 41
Chapter Problems 43
CHAPTER 2

Elements and Compounds
PA R T

A

50

THE ELEMENTS AND THEIR COMPOSITION

2-1 The Elements

52

52

MAKING IT REAL:

Iridium, the Missing Dinosaurs, and the Scientific Method

55

2-2 The Composition of Elements: Atomic Theory 56
2-3 Composition of the Atom 58
2-4 Atomic Number, Mass Number, and Atomic Mass 60
MAKING IT REAL:

Isotopes and the History of Earth’s Weather

PA R T A S U M M A R Y

62

64

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CONTENTS

PA R T

B

COMPOUNDS AND THEIR
COMPOSITION
64

2-5 Molecular Compounds
2-6 Ionic Compounds 67
MAKING IT REAL:

64
Ionic Compounds and Essential Elements

PA R T B S U M M A R Y

Chapter Summary
Chapter Problems

71

71

72
74

CHAPTER 3

The Properties of Matter and Energy
PA R T

A

T H E P R O P E R T I E S O F M AT T E R

3-1 The Physical and Chemical Properties of Matter
3-2 Density—A Physical Property 84
MAKING IT REAL:

PA R T A S U M M A R Y

B

80

80

Identifying a Glass Shard from a Crime Scene by Density

3-3 The Properties of Mixtures

PA R T

78

89
93

TH E PR OPE RTI ES OF E N E R GY

3-4 The Forms and Types of Energy 94
3-5 Energy Measurement and Specific Heat

94

96

MAKING IT REAL:

Body Solutions—Lose Weight (actually money)
While You Sleep
99

MAKING IT REAL:

Dark Matter and Energy

PA R T B S U M M A R Y

Chapter Summary
Chapter Problems

88

100

101

102
104

CHAPTER 4

The Periodic Table and Chemical Nomenclature
PA R T

A

R E L AT I O N S H I P S A M O N G T H E E L E M E N T S
A N D T H E P E R I O D I C TA B L E
110

4-1 The Origin of the Periodic Table
4-2 Using the Periodic Table 112
MAKING IT REAL:

B

110

The Discovery of a Group VIIA Element, Iodine

PA R T A S U M M A R Y

PA R T

108

114

116

THE FORMULAS AND NAMES
OF COMPOUNDS
116

4-3 Naming and Writing Formulas of Metal-Nonmetal Binary Compounds 117
4-4 Naming and Writing Formulas of Compounds with Polyatomic Ions 121
M A K I N G I T R E A L : Ionic Compounds in the Treatment of Disease
124
4-5 Naming Nonmetal-Nonmetal Binary Compounds 125
4-6 Naming Acids 127
PA R T B S U M M A R Y
129
Chapter Summary 130
Chapter Problems 132

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CONTENTS

CHAPTER 5

Chemical Reactions
PA R T

A

136

T H E R E P R E S E N TAT I O N O F C H E M I C A L
CHANGES AND THREE TYPES
OF CHANGES
138

5-1 Chemical Equations 138
5-2 Combustion, Combination, and Decomposition Reactions
MAKING IT REAL:

PA R T A S U M M A R Y

PA R T

B

143

Life Where the Sun Doesn’t Shine—Chemosynthesis

I O N S I N W AT E R A N D H O W T H E Y R E A C T

5-3 The Formation of Ions in Water 148
5-4 Single-Replacement Reactions 150
5-5 Double-Replacement Reactions–Precipitation
MAKING IT REAL:

146

147

153

Hard Water and Water Treatment

5-6 Double-Replacement Reactions–Neutralization
PA R T B S U M M A R Y
161
Chapter Summary 162
Chapter Problems 164

1 47

158

159

CHAPTER 6

Quantities in Chemistry
PA R T

A

168

THE MEASUREMENT OF MASSES OF
ELEMENTS AND COMPOUNDS
170

6-1 Relative Masses of Elements 170
6-2 The Mole and the Molar Mass of Elements
6-3 The Molar Mass of Compounds 177
PA R T A S U M M A R Y
181
PA R T

B

173

THE COMPONENT ELEMENTS OF
COMPOUNDS
182

6-4 The Composition of Compounds
MAKING IT REAL:

182

Calcium in Our Diets—How Much Are You Getting?

6-5 Empirical and Molecular Formulas
PA R T B S U M M A R Y
190
Chapter Summary 191
Chapter Problems 192

186

187

CHAPTER 7

Quantitative Relationships in Chemical
Reactions
196
PA R T

A

M A S S R E L AT I O N S H I P S I N C H E M I C A L
REACTIONS
198

7-1 Stoichiometry 198
7-2 Limiting Reactant 203
MAKING IT REAL:

7-3 Percent Yield

Alcohol—The Limiting Reactant in Breathalyzers

208

PA R T A S U M M A R Y

210

207

xi

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CONTENTS

PA R T

B

E N E R G Y R E L AT I O N S H I P S I N C H E M I C A L
REACTIONS
211

7-4 Heat Energy in Chemical Reactions
MAKING IT REAL:

213

214

PA R T B S U M M A R Y

Chapter Summary
Chapter Problems

211

Hydrogen—The Perfect Fuel

214
215

CHAPTER 8

Modern Atomic Theory
PA R T

A

220

TH E E N E R GY OF TH E E LECTR ON
I N T H E AT O M
222

8-1 The Emission Spectra of the Elements and Bohr’s Model
MAKING IT REAL:

M A K I N G I T R E A L : Solving Crimes with Light
227
8-2 Modern Atomic Theory: A Closer Look at Energy Levels
PA R T A S U M M A R Y
232

PA R T

B

222

Roses Are Red; Violets Are Blue—But Why?

223

227

T H E P E R I O D I C TA B L E A N D E L E C T R O N
C O N F I G U R AT I O N
233

8-3 Electron Configurations of the Elements 233
8-4 Orbital Diagrams of the Elements (Optional) 240
8-5 Periodic Trends 242
PA R T B S U M M A R Y
245
Chapter Summary 246
Chapter Problems 248

CHAPTER 9

The Chemical Bond
PA R T

A

252

C H E M I C A L B O N D S A N D T H E N AT U R E
OF IONIC COMPOUNDS
254

9-1 Bond Formation and Representative Elements
9-2 Formation of Ions and Ionic Compounds 255
PA R T A S U M M A R Y
259
PA R T

B

254

C H E M I C A L B O N D S A N D T H E N AT U R E
OF MOLECULAR COMPOUNDS
260

9-3 The Covalent Bond

260

MAKING IT REAL:

Nitrogen: From the Air to Proteins

263

9-4 Writing Lewis Structures 264
9-5 Resonance Structures 269
PA R T B S U M M A R Y
271
PA R T

C

THE DISTRIBUTION OF CHARGE
IN CHEMICAL BONDS
272

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CONTENTS

9-6
9-7
9-8
9-9

Electronegativity and Polarity of Bonds
Geometry of Simple Molecules 275
Polarity of Molecules 278
Formal Charge (Optional) 280
MAKING IT REAL:

Enzymes—The Keys of Life

281

284

PA R T C S U M M A R Y

Chapter Summary
Chapter Problems

272

285
287

CHAPTER 10

The Gaseous State
PA R T

A

292

T H E N AT U R E O F T H E G A S E O U S
S TAT E A N D T H E E F F E C T S O F
CONDITIONS
294

10-1 The Nature of Gases and the Kinetic Molecular Theory
M A K I N G I T R E A L : Ozone—Friend and Foe
296
10-2 The Pressure of a Gas 297
10-3 Charles’s, Gay-Lussac’s, and Avogadro’s Laws 301
PA R T A S U M M A R Y
309
PA R T

B

294

R E L AT I O N S H I P S A M O N G Q U A N T I T I E S
OF GASES, CONDITIONS, AND CHEMICAL
REACTIONS
309

10-4 The Ideal Gas Law 310
10-5 Dalton’s Law of Partial Pressures 313
10-6 The Molar Volume and Density of a Gas
MAKING IT REAL:

316

Defying Gravity—Hot-Air Balloons

319

10-7 Stoichiometry Involving Gases 319
PA R T B S U M M A R Y
321
Chapter Summary 322
Chapter Problems 324

CHAPTER 11

The Solid and Liquid States
PA R T

A

330

THE PROPERTIES OF CONDENSED
S TAT E S A N D T H E F O R C E S
I N V O LV E D
332

11-1 Properties of the Solid and Liquid States 332
11-2 Intermolecular Forces and Physical State 334
11-3 The Solid State: Melting Point 338
MAKING IT REAL:

The Melting Point of Iron and the World Trade
Center
341

PA R T A S U M M A R Y

342

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CONTENTS

PA R T

B

T H E L I Q U I D S TAT E A N D C H A N G E S
I N S TAT E
343

11-4 The Liquid State: Surface Tension, Viscosity,
and Boiling Point 343
MAKING IT REAL:

The Oceans of Mars

348

11-5 Energy and Changes in State 349
11-6 Heating Curve of Water 352
PA R T B S U M M A R Y
355
Chapter Summary 356
Chapter Problems 358

CHAPTER 12

Aqueous Solutions
PA R T

A

362

SOLUTIONS AND THE QUANTITIES
I N V O LV E D
364

12-1 The Nature of Aqueous Solutions 364
12-2 The Effects of Temperature and Pressure
on Solubility 367
MAKING IT REAL:

Hyperbaric Therapy

369

12-3 Concentration: Percent by Mass 370
12-4 Concentration: Molarity 373
12-5 Stoichiometry Involving Solutions 377
PA R T A S U M M A R Y
381
PA R T

B

THE EFFECTS OF SOLUTES ON THE
P R O P E R T I E S O F W AT E R
382

12-6 Electrical Properties of Solutions 382
12-7 Colligative Properties of Solutions 384
MAKING IT REAL:

Osmosis in a Diaper

PA R T B S U M M A R Y

Chapter Summary
Chapter Problems

389

390

390
392

CHAPTER 13

Acids, Bases, and Salts
PA R T

13-1
13-2
13-3
13-4

A

396

A C I D S , B A S E S , A N D T H E F O R M AT I O N
O F S A LT S
398

Properties of Acids and Bases 398
Brønsted–Lowry Acids and Bases 400
Strengths of Acids and Bases 404
Neutralization and the Formation of Salts
MAKING IT REAL:

408

Salts and Fingerprint Imaging

PA R T A S U M M A R Y

412

411

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CONTENTS

PA R T

B

THE MEASUREMENT OF ACID
STRENGTH
412

13-5 Equilibrium of Water
13-6 The pH Scale 415

412

PA R T B S U M M A R Y

PA R T

C

419

S A LT S A N D O X I D E S A S A C I D S
AND BASES 420

13-7 The Effect of Salts on pH—Hydrolysis 420
13-8 Control of pH—Buffer Solutions 423
MAKING IT REAL:

The pH Balance in the Blood

13-9 Oxides as Acids and Bases
MAKING IT REAL:

426

Acid Rain—The Price of Progress?

PA R T C S U M M A R Y

Chapter Summary
Chapter Problems

425
428

428

429
431

CHAPTER 14

Oxidation–Reduction Reactions
PA R T

A

436

R E DOX R EACTIONS—TH E EXCHANG E
438
OF ELECTRONS

14-1 The Nature of Oxidation and Reduction and Oxidation States
MAKING IT REAL:

438

Lightning Bugs (Fireflies)—Nature’s Little Night-Lights

14-2 Balancing Redox Equations: Oxidation State Method 444
14-3 Balancing Redox Equations: Ion-Electron Method 447
PA R T A S U M M A R Y
451
PA R T

B

S P O N TA N E O U S A N D N O N S P O N TA N E O U S
R E DOX R EACTIONS
452

14-4 Predicting Spontaneous Redox Reactions
14-5 Voltaic Cells 457
MAKING IT REAL:

14-6 Electrolytic Cells

453

Fuel Cells—The Future Is (Almost) Here

462

PA R T B S U M M A R Y

Chapter Summary
Chapter Problems

464

465
466

CHAPTER 15

Reaction Rates and Equilibrium
PA R T

A

470

COLLISIONS OF MOLECULES AND
R E A C T I O N S AT E Q U I L I B R I U M
472

15-1 How Reactions Take Place 472
15-2 Rates of Chemical Reactions 475
15-3 Equilibrium and Le Châtelier’s Principle
MAKING IT REAL:

479

The Lake That Exploded

PA R T A S U M M A R Y

484

483

462

443

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CONTENTS

PA R T

B

T H E Q U A N T I TAT I V E A S P E C T S O F
R E A C T I O N S AT E Q U I L I B R I U M
485

15-4 The Equilibrium Constant 485
15-5 Equilibria of Weak Acids and Weak Bases in Water
MAKING IT REAL:

15-6 Solubility Equilibria

496

497

PA R T B S U M M A R Y

Chapter Summary
Chapter Problems

489

Buffers and Swimming Pool Chemistry

502

502
505

CHAPTER 16

Nuclear Chemistry
PA R T

A

510

N A T U R A L LY O C C U R R I N G
RADIOACTIVITY
512

16-1 Radioactivity 512
16-2 Rates of Decay of Radioactive Isotopes
16-3 The Effects of Radiation 518
MAKING IT REAL:

516

Radioactivity and Smoke Detectors

16-4 The Detection and Measurement of Radiation
PA R T A S U M M A R Y
523
PA R T

B

520

521

INDUCED NUCLEAR CHANGES
AND THEIR USES
524

16-5 Nuclear Reactions 524
16-6 Applications of Radioactivity 526
16-7 Nuclear Fission and Fusion 529
MAKING IT REAL:

Revisiting the Origin of the Elements

PA R T B S U M M A R Y

Chapter Summary
Chapter Problems

534

535
537

C H A P T E R 17 ( W E B )

Organic Chemistry
PA R T

A

w-2

HYDROCARBONS

17-1 Bonding in Organic Compounds
17-2 Alkanes w-8
17-3 Alkenes and Alkynes w-13
MAKING IT REAL:

17-4 Aromatic Compounds
PA R T A

w-4

w-4

The Age of Plastics

w-17
SUMMARY
w-19

w-16

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CONTENTS

PA R T

17-5
17-6
17-7
17-8

B

OTHER CLASSES OF ORGANIC
COMPOUNDS
W-20

Alcohols and Ethers w-20
Aldehydes and Ketones w-22
Amines w-24
Carboxylic Acids, Esters, and Amides
MAKING IT REAL:

w-28

w-29

PA R T B S U M M A R Y

Chapter Summary
Chapter Problems

w-25

Aspirin—An Old Drug with a New Life

w-30
w-32

CHAPTER 18 (WEB)

Biochemistry
PA R T

A

w-42

THREE BASIC TYPES OF BIOCHEMICAL
COMPOUNDS
w-44

18-1 Lipids

w-44

MAKING IT REAL:

Soap—Dissolves Away Fat

18-2 Carbohydrates w-49
18-3 Amino Acids and Proteins

w-53
w-57

PA R T A S U M M A R Y

PA R T

B

w-48

BIOCHEMICAL COMPOUNDS AND LIFE
FUNCTIONS
W-58

18-4 Enzymes w-58
18-5 Nucleic Acids and Genetics

w-60

The Origin of Life

MAKING IT REAL:

w-65

PA R T B S U M M A R Y

Chapter Summary
Chapter Problems

F O R E W O R D

TO

w-64

w-66
w-68

T H E

A P P E N D I C E S

APPENDIX A:

B A S I C M AT H E M AT I C S

APPENDIX B:

BASIC ALG E B R A

APPENDIX C:

S C I E N T I F I C N O TAT I O N

APPENDIX D:

GRAPHS

APPENDIX E:

C A L C U L AT O R S

APPENDIX F:

G LOSSARY

A-10
A-19

A-27
A-31

A-36

A P P E N D I X G : A N S W E R S TO C H A P T E R P R O B L E M S
P H OTO C R E D I T S
INDEX

A-1

A-2

I-1

PC-1

A-45

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P R E FA C E

Leo Malone, sole author of the first seven editions, has joined forces with Ted Dolter
to create Basic Concepts of Chemistry, Eighth Edition. Professor Dolter’s academic
background in chemical education and professional duties as chairman of the
chemistry program at Southwestern Illinois College uniquely qualifies him to integrate
this new dimension of outcomes assessment into Basic Concepts of Chemistry.
Although the new edition continues its focus on the preparatory and basic chemistry
market, it now includes emphasis on chapter objectives and their assessment.

Why Did We Write This Book?
Basic Concepts of Chemistry was originally written in 1981 to address the needs of
students planning to take the general chemistry course, but with little or no
background in chemistry. Over the next seven editions, its mission has evolved, so
that in this new eighth edition we have focused on integrating meaningful
assessment and timely feedback into the book as well as the accompanying online
course management course (WileyPLUS). This book has been used extensively in
one semester, general-purpose courses, where professors find students at a variety
of levels. For some of these students a main sequence in chemistry may follow, but
for others the course precedes a semester of organic and biochemistry. Still others
enroll to satisfy a science requirement or a one semester stand-alone chemistry
requirement. The text was written at a level and with the functionality designed to
accommodate the needs of each of these groups of students, by structuring itself so
that an instructor could emphasize or omit certain clearly delineated sections.

Basic Concepts of Chemistry Today
Today, more and more students are entering post secondary education with a wider
variety of learning styles and varied levels of preparedness. Recent data on student
populations indicate that many of them are visual and kinesthetic learners and it is
vital that textbook authors both recognize and integrate this pedagogy into their
books. We have focused hard on accomplishing this goal. Furthermore, the
students that take this course have very diverse math backgrounds and we have
provided many valuable resources to allow those students with weak math skills to
succeed in this course. To accommodate the visual and kinesthetic learner and to
support those students with weaker math backgrounds, we have introduced several
new features, principle among these is a focus on outcomes assessment.

The Role of Outcomes Assessment in this Edition
Surveys of programs across the country show that chemistry is being taught in
numerous ways. From packed lecture halls to intimate classrooms, from Ph.D.s to
adjunct instructors to teaching assistants, chemistry education is being delivered in
multiple formats, often within the same institution. Outcomes Assessment is an

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P R E FA C E

attempt to insure consistency in evaluating student achievement across these
multiple formats. By delineating the expected outcomes or objectives for each
chapter, and then devising assessment tools, such as homework and exam
questions, lab experiments, group work, and the like, that are designed to target
those specific objectives, schools can insure that all students in all sections are being
served. By incorporating outcomes assessment into the curriculum, students can
receive the same topical instruction and be evaluated against the same standard.
Implementing and justifying an outcomes assessment program can be time
consuming. To that end, the authors, having already been through an outcomes
assessment program, have designed a text with objectives and assessments already
in place, along with the required information needed to show how all of it ties
together. Each Part within chapters includes a list of the relevant objectives for that
chapter.

PA R T A
THE PROPERTIES
O F M AT T E R

OBJ ECTIVES

SETTI NG A GOAL
■

You will learn how a sample of matter can be described
by its properties and how they can be quantitatively
expressed.

3-1 List and define several properties of matter and
distinguish them as physical or chemical.
3-2 Perform calculations involving the density of liquids
and solids.
3-3 (a) Describe the differences in properties between a
pure substance and a mixture. (b) Perform calculations
involving percent as applied to mixtures.

These are measurable outcomes that the student should master by the completion
of that part. Assessments of varying complexity follow each section so that the student,
upon completion, can evaluate to what degree the material has been internalized.

4 - 3 ( a ) K N O W L E D G E : For which of the following metals
must a charge be placed in parenthesis when naming one of its compounds?
Co, Li, Sn, Al, Ba

EXERCISE

E X E R C I S E 4 - 3 ( b ) A N A LY S I S :

(a) Li2O

(b) CrI3

(c) PbS

ASSESSING THE
OBJECTIVE FOR
SECTION 4-3

Name the following compounds.
(d) Mg3N2
(e) Ni3P2

Provide the formula for the following chemicals.
(c) tin(IV) bromide
(d) calcium nitride

E X E R C I S E 4 - 3 ( c ) A N A LY S I S :

(a) aluminum iodide
(b) iron(III) oxide

Using an M to represent the metal and an
X to represent the nonmetal, write theoretical formulas for all possible combinations of M and X with charges of 1, 2, 3 and 1, 2, and 3.

EXERCISE 4-3(d) S Y N T H E S I S :

For additional practice, work chapter problems 4-18, 4-20, 4-22, and 4-24.

At the end of each chapter summary, there is an objectives grid which ties the
objectives to examples within the sections, assessment exercises at the end each
section, and relevant chapter problems at the end of each chapter. This grid has
utility for both the student and instructor so that both can easily locate tools within
the text to help address a specific student problem.

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P R E FA C E

OBJECTIVES
SECTION YOU SHOU LD B E AB LE TO...

EXAM PLES

EXE RCISES

CHAPTE R PROB LE MS

5-1

Write a chemical equation from a word description
of the reaction.

5-1

1a, 1b, 1e, 2c

10, 11, 12, 13

Balance a simple chemical equation by inspection.

5-1, 5-2, 5-3

1c, 1d

2, 3, 4, 5

5-2

Classify certain chemical reactions as being
combustion, combination, or decomposition reactions.

5-3

2a, 2b, 2c

14, 15, 16, 18, 19,
20, 21, 22, 23

5-3

Write the ions formed when ionic compounds or
acids dissolve in water.

5-4

3a, 3b, 3c

24, 25, 26, 27

5-4

Given the activity series, complete several
single-replacement reactions as balanced molecular,
total ionic, and net ionic equations.

5-5

4a, 4b, 4c, 4d

28, 29, 30, 33

5-5

Given a table of solubility rules, determine whether
a specific ionic compound is soluble or insoluble
in water.

5-6

5a, 5b, 5c

34, 35, 40, 42

Write balanced molecular, total ionic, and net ionic
equations for precipitation reactions.

5-7, 5-8, 5-9

5d, 5e

44, 45, 46, 47,
48, 49

Write balanced molecular, total ionic, and net ionic
equations for neutralization reactions.

5-10

6a, 6b

55, 56, 57, 58

5-6

Other changes to the eighth edition include:
• Emphasizing outcomes assessment, each Chapter Part now begins with Setting
Goals, which serves to preview the important topics within that Part. Section
Objectives are presented alongside these Goals so that students can realistically
link the text’s goals to the objectives required of them.
• Assessing the Objectives are a collection of problems that appear at the end of
each section, and provide students with the opportunity to assess their
understanding of each section’s objectives. The representative problems are
divided into three cognitive levels: Knowledge, Analysis, and Synthesis. Each level is
progressively more sophisticated and prompts the students to gauge their
conceptual understanding of the section content.
• Every concept in the text is clearly illustrated with one or more step by step
examples. Most examples, in addition to a Procedure and Solution step, are
followed by two new steps Analysis and Synthesis. The Analysis step discusses the
problem in light of the reasonableness of the answer, or perhaps suggests an
alternate way to solve the problem involving different learning modes. The
concluding Synthesis step gives the student the opportunity to delve deeper,
asking the student to extend their knowledge. These added steps promote
critical thinking and facilitate deeper conceptual understanding.
• Making it Real essays have been updated to present timely and engaging realworld applications, emphasizing to the student the relevance of the material they
are learning. For example, in the high-interest field of forensics we describe how
glass shards from crime scenes can be identified by their density in Chapter 3,
and then in Chapter 8, explore how the refractive index of glass is also used as
important evidence in solving crimes. In Chapter 7, we take a look at the
chemical reactions involved in the breathalyzer and in Chapter 13, how salts are
used to analyze fingerprints.
• New to this edition are end of chapter Student Workshop activities. These are
intended to cater to the many different student learning styles and to engage them
in the practical aspect of the material discussed in the chapter. Each “Student
Workshop” includes a statement of purpose and an estimated time for completion.

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P R E FA C E

Organization
The Prologue is a unique feature which introduces the origin of science in general
and chemistry in particular. There are no quizzes, exercises, or problems; rather, it
is meant to be a relaxing, historical glimpse at the origin of this fascinating subject
and how it now affects our lives. Our intent is build interest and engage the student
in further study. We recognize the changing needs of students and balance that with
the requirements to successfully study chemistry. As such, we continue to provide
the necessary support for students continuing on in the study of chemistry. Chapter
1, Measurements in Chemistry, provides the necessary math tools in a nonthreatening way. In Chapter 2, Elements and Compounds, the elements are
introduced starting from what we see and sense about us (the macroscopic) to the
atoms of which they are composed and finally into the structure within the atom
(the microscopic and submicroscopic.) We do the same for compounds in the
second part of this chapter. Chapter 3, The Properties of Energy and Matter
continues the discussion of matter and its properties. Some additional yet relevant
math concepts such as density, percent composition, and specific heat are
introduced in this chapter as properties of matter. Chapter 4, The Periodic Table
and Chemical Nomenclature, allows us to draw in one of the primary tools of the
chemist, the periodic table. We see its functionality and organization, and begin
using it in a thorough discussion of how to name most common chemicals, whose
structure was discussed in chapter 2.
In Chapter 5, Chemical Reactions we discuss the broad range of chemical
interactions that can occur. The quantitative aspects of chemistry are discussed in
Chapter 6, Quantities in Chemistry and Chapter 7, Quantitative Relationships in
Chemical Reactions. These chapters were split from a single, larger chapter in the
7th edition, and have been moved forward ahead of Chapter 8, Modern Atomic
Theory and Chapter 9, The Chemical Bond. Still, the two latter chapters can be
moved ahead of Chapter 5 without prejudice, depending on the preferences of the
instructor, and the ease with which the material can be incorporated into the
overall curriculum.
Chapter 10, The Gaseous State, begins a three chapter in depth study of the
states of matter by examining the unique and predictable behaviors of gases. This
is followed by similar discussions of the condensed states of matter in Chapter 11,
The Liquid and Solid States, which includes a thorough but understandable
discussion of intermolecular forces and how those affect the properties of matter.
This discussion continues in Chapter 12, Aqueous Solutions, where we discuss
both the qualitative and quantitative aspects of solute-solvent interactions. This
chapter serves as a wrap-up of the quantitative relationships introduced in
Chapters 6 and 7.
Chapter 13, Acids, Bases and Salts, begins a discussion of specific classes of
chemicals, and goes into depth as regards the ways acids and bases can react
together and with their environment. A second class of reaction is explored in
depth in Chapter 14, Oxidation-Reduction Reactions. The relationship of these
types of reactions to the creation of batteries and electrical currents is explained.
The remaining chapters offer a survey of topics that are of general chemical
interest, and are appropriate for those looking to expose their students to a broad
set of topics. Chapter 15, Reaction Rate and Equilibrium, introduces the concepts
of Kinetics and Equilibrium and gives a taste of the sophisticated mathematical
treatment these topics receive. Chapter 16, Nuclear Chemistry, explores how
radioactivity is a phenomenon associated with certain elements and isotopes.
Chapter 17, Organic Chemistry, gives the briefest introduction to organic
functional groups, structure, and bonding. Chapter 18, Biochemistry, gives a
similar treatment to common biochemical structures like carbohydrates, proteins,
lipids, and nucleic acids.

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P R E FA C E

Supplements
WileyPLUS with CATALYST
WileyPLUS is a powerful online tool that provides a completely integrated suite of
teaching and learning resources on one easy to use website. WileyPLUS integrates
Wiley’s world-renowned content with media, including a multimedia version of the
text, PowerPoint slides, digital image archive, online assessment, and more.
WileyPLUS with CATALYST partners with the instructor to teach students how
to think their way through problems, rather than rely on a list of memorized
equations, by placing a strong emphasis on developing problem solving skills and
conceptual understanding. WileyPLUS with CATALYST incorporates an online
learning system designed to facilitate dynamic learning and retention of learned
concepts. CATALYST was developed by Dr. Patrick Wegner (California State
University, Fullerton) to promote conceptual understanding and visualization of
chemical phenomena in undergraduate chemistry courses.
CATALYST assignments have multiple levels of parameterization and test on
key concepts from multiple points of view (visual, symbolic, graphical,
quantitative). Hundreds of end-of-chapter problems are available for assignment,
and all are available with multiple forms of problemsolving support.
Study Guide/Solutions Manual by Leo J. Malone is available to accompany this text. In
the Study Guide/Solutions Manual, the same topics in a specific section are also
grouped in the same manner for review, discussion, and testing. In this manner, the
Study Guide/Solutions Manual can be put to use before the chapter is completed.
The Study Guide/Solutions Manual contains answers and worked-out solutions to
all problems in green lettering in the text.
Experiments in Basic Chemistry by Steven Murov and Brian Stedjee, Modesto Junior
College. Taking an exploratory approach to chemistry, this hands-on lab manual for
preparatory chemistry encourages critical thinking and allows students to make
discoveries as they experiment. The manual contains 26 experiments that parallel
text organization and provides learning objectives, discussion sections outlining each
experiment, easy-to-follow procedures, post-lab questions, and additional exercises.
Instructor’s Manual and Test Bank by Leo J. Malone, St Louis University; Ted Dolter and
Steve Gentemann, Southwestern Illinois College; and Kyle Beran, University of TexasPermian Basin. The Instructor’s Manual consists of two parts; the first part includes
hints and comments for each chapter by Ted Dolter, followed by daily lesson plans by
Steve Gentemann. The second part of the manual contains answers and worked-out
solutions to all chapter-end problems in the text by Leo Malone. The test bank by
Kyle Beran consists of multiple choice, short answer, and fill in the blank questions.
Instructor’s Manual for Experiments in Basic Chemistry, written by the lab manual
authors, contains answers to post-lab questions, lists of chemicals needed,
suggestions for other experiments, as well as suggestions for experimental set-ups.
Power Point Lecture Slides, created by Wyatt Murphy, of Seton Hall University, these
slides contain key topics from each chapter of the text, along with supporting
artwork and figures from the text. The slides also contain assessment questions that
can be used to facilitate discussions during lecture.
Digital Image Archive. The text web site includes downloadable files of text images
in JPEG format.
Computerized Test Bank The IBM and Macintosh compatible version of the entire
Test Bank has full editing features to help the instructor customize tests.
Most of the resources listed above can also be accessed at www.wiley.com/
college/malone.

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A C K N O W L E D G M E N T S

A revision of this magnitude involves efforts spanning several years and requiring
the input of many people. In particular, Dr. Malone thanks his colleagues at Saint
Louis University for their helpful comments in previous editions of this text. He’s
grateful to his wife Meg, who demonstrated patience and put up with occasional
crabbiness during the new text’s preparation. Dr. Malone also appreciates the
support of his children and their spouses: Lisa and Chris, Mary and Brian, Katie
and Rob, and Bill. They and their eleven children were both a source of great
inspiration and a large amount of noise.
Professor Dolter thanks his general chemistry instructor, who provided a sound
understanding of the basic principles needed in forming his craft. Although his
professor (LJM) was a taskmaster, the end result was worth it. He’d also like to
thank his colleagues at SWIC, who happily allowed him to bounce ideas off of them,
and delighted in giving their advice. Professor Dolter is thankful for his wife Peggy,
who endured the single-parent lifestyle many weekends during the writing of this
text. A nod also goes to, Isabel and Zachary, who wondered if their dad was ever
going to finish this book; they’re owed some dad-time.
The authors also wish to thank the many people at John Wiley who helped and
encouraged this project, especially Joan Kalkut for keeping us moving. In addition,
we’d like to thank Nick Ferrari, our editor, for his support of this revision; Jennifer
Yee, our supplements editor; Cathy Donovan, chemistry assistant extraordinaire;
photo researcher, Lisa Gee; and Suzanne Ingrao of Ingrao Associates, who kept an
eye on every production detail, and kindly made sure we kept on schedule. Kevin
Murphy and Brian Salisbury did a wonderful job of creating the new design, which
highlights the Outcomes Assessment focus of the text. Thank you all for remaining
wonderfully patient in the face of missed deadlines and family conflicts.
Finally, the following people offered many useful comments and suggestions for the
development of the Eighth Edition:
Jeanne C. Arquette, Phoenix College
Gerald Berkowitz, Erie Community College
Ken Capps, Central Florida Community College
Brant Chapman, Bergen Community College
Douglas S. Cody, Nassau Community College
John J. Dolhun, Norwalk Community College
Danae Quirk Dorr, Minnesota State University, Mankato
Rob Fremland, San Diego Mesa College
Amy Grant, El Camino College
Paul R. Haberstroh, Mohave Community College
Mary Hadley, Minnesota State University
Alton Hassell, Baylor University
Bettina Heinz, Palomar College
Bruce E. Hodson, Baylor University

Tracy Holbrook, Cape Fear Community College
Larry Kolopajlo, Eastern Michigan University
Virginia R. Kotlarz, Erie Community College
Rosemary Leary, Estrella Community College
Rich Lomneth, University of Nebraska at Omaha
Kathy Mitchell, St. Petersburg College
Krzysztof Ochwat, Wilbur Wright College
Neil Palosaari, Inver Hills Community College
Todd Rogers, Columbia Basin College
Leo J. Malone
Saint Louis University
Theodore O. Dolter
Southwestern Illinois College

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BASIC CONCE PTS
OF CH E M ISTRY

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P R O L O G U E

Science and
the Magnificent
Human Mind
hat could be more peaceful than a campfire silhouetted
by the night sky? Actually, the sky and the fire were the
source of remarkable achievements of the human race but in
different ways. Observations of the sky and the stars provided
the keys, in an ongoing process, that are unlocking the
secrets of space, time, and the history of the cosmos.
The use of fire, an awesome force of nature,
provided the protection and the warmth that
allowed the human race to thrive in a
hostile world. Eventually, the use of fire
would be the instrument of huge
advances in civilization. The
Prologue relates how the study
of the stars and the use of
fire led us to the world of
science that serves us so
ably today.

W

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B
C

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The Origin of Matter
The Mystery of Fire
The Scientific Method

Page 3

A M O N G T H E A N I M A L K I N G D O M , only
humans have the ability to take their minds beyond simple
survival. We also analyze, ponder, and predict the future
based on observations. This has led us to a remarkable
understanding of all that we see and otherwise sense about
us. So this Prologue is dedicated to how the wonderful
workings of our minds have allowed us to establish the
realm of modern science. In Section A, we examine how the
first stirrings of curiosity about the night sky led eventually
to the still developing understanding of what has happened
since the beginning of time. As presented here, we are preempting the use of some terms and concepts that will be
explained more thoroughly in later chapters. The readings in
the Prologue are meant to tweak your interest, not necessarily to be learned. In Section B, we see how the taming of fire
by our ancient ancestors propelled us into huge scientific
discoveries and the basis of chemistry. Finally, in Section C,
we take note of how science in general and chemistry in
particular have progressed from random discoveries and
serendipity to the complex technical world in which we exist
today. After proceeding through the Prologue, we will build
the topic of chemistry from the most basic substances of the
universe to the complex world of chemistry that serves us in
so many ways today.

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PROLOGUE

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Science and the Magnificent Human Mind

A

T H E O R I G I N O F M AT T E R

Somewhere in ancient time, many thousands of years ago, someone looked up at
the night sky and wondered what this display of bright dots meant. It was a great
landmark in our existence, as humans became able to expand their minds beyond
just the concerns of everyday survival and existence. With the power of abstract reasoning, our earliest ancestors began to observe, consider, and speculate about the
world around them. Surely the glittering display of bright stars in the heavens would
have attracted the attention of their newly inquisitive minds. One can only imagine
what they must have thought they were observing. They probably sensed that something very vast and mysterious was involved, but they would not have known that the
secrets of our very existence are indeed “written in the stars.” What cosmologists have
garnered from the night sky in just the past few decades has given us a tantalizing
yet still incomplete understanding of the beginning, evolution, and ultimate destiny
of our entire universe.
From observations of the cosmos, many of the mysteries of the origin of all that
we see, from the paper in front of us to the farthest reach of the cosmos in infinite space, have been revealed. On a small piece of dust in the universe, called
Earth, we peer out into distant space and back into ancient time. We observe the
flickering points of light that we call stars—the visible flashes from distant suns,
some like our own. Magnification of the heavens tells us that some of what appear
as individual stars to the naked eye are actually groups of stars called galaxies, each
containing hundreds of billions of individual stars. It was the motions of these
galaxies in the heavens that provided modern scientists with the clues about the
origin of everything. One observation divulged that groups of galaxies seem to be
pushing away from each other, indicating a dynamic, continually expanding universe. A logical conclusion from this is that the galaxies were closer together in
the past. In fact, if we go back a little less than 14 billion years, the entire cosmos
was coalesced into a single, infinitely dense point known as the singularity. The
spontaneous expansion of this point (which is referred to as “the big bang”)
marked the beginning of time and space. At the very start of time, the two components of the universe—the stars, planets, and other forms of matter (which has
mass) as well as the heat, light, and other forms of energy (which has no mass)—
were to evolve from the singularity. How the singularity came to be and what was
present before time began is unknown to science and may never be known. But,
immediately after the big bang, there was only energy and unbelievable heat. It
wasn’t for another 10,000 years that our infant universe expanded and cooled
enough so that some of the energy of this nascent system evolved into the basic
building blocks of nature.
The first matter was composed of three basic particles known as protons, neutrons, and electrons, which evolved from energy and even more fundamental particles. At first they existed independently in the hot, dense universe. These
particles exist in the world of the submicroscopic—they are way too small to see
with even the most powerful magnification. The proton carries a positive electrical charge; the neutron, about the same mass as a proton, carries no charge. The
third particle, the electron, tiny compared to protons and neutrons (about 1/2000
the mass), carries a negative charge equal but opposite to the charge on the proton. From these three particles, the entire stuff of the universe would eventually
evolve.
About 400,000 years after the big bang, a hugely significant event occurred. The
particles slowed down enough so that the attractive forces between a positively
charged proton and a negatively charged electron could hold the two particles

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A The Origin of Matter

together. Significantly, the electron and proton did not directly attach to each other.
Instead, the electron settled into a stable orbit around the central proton. Why this
happened is not easily explained but is understood in the realm of quantum mechanics. Quantum mechanics is the science that deals with the forces involved in the very
small dimensions of these particles. For now we can accept some results of more complex theories without going into detail.
A model of the hydrogen atom, with an electron in a set orbit, was first advanced
by Niels Bohr in 1923. It compares the orbiting of the electron around the proton
to the motion of the Earth around the sun. In fact, the nature of the electron in the
atom is now known to be more complex, but this classical orbiting model serves us
well at this point, so we will use it. The proton and the electron together now formed
a neutral atom. Normal matter is composed of variations of these tiny particles called
atoms. The smallness of atoms is a stretch for the human mind to comprehend. In
fact, the period at the end of this sentence contains a number of atoms comparable to the number of grains of sand in the Sahara Desert (about 1017, or one hundred million billion atoms).
The proton is the center of this tiny two-particle universe and is known as the
nucleus of the atom. The vast volume of the atom is empty space in which the electron exists. If the minute size of the atom isn’t hard enough to imagine, the dimensions within the atom itself are equally difficult for our modest and limited minds
to fathom. Imagine that the nucleus was the size of a basketball. In that case, the
electron, which defines the total size of the atom, would exist in a volume with a
two-mile radius. This model represents an atom of the element hydrogen, which
accounted for 90% of the atoms initially formed in the big bang. An element is a basic
form of matter.

–

Hydrogen
atom

+

What of the other 10% of the matter that was formed? Some of the nuclei of
hydrogen contained one neutron as well as one proton. This form of hydrogen is
often referred to as deuterium or heavy hydrogen, but it is not a unique element. About
one out of 30,000 hydrogen atoms are actually deuterium atoms. Atoms of a certain
element that have a specific number of neutrons are known as isotopes. So, initially,
the element hydrogen existed as a mixture of two isotopes, normal hydrogen and a
very small amount of deuterium.
The earliest universe contained another type of nucleus that had two protons and
two neutrons. When combined with two electrons, these nuclei formed the atoms
of a second element known as helium. The atoms of a specific element have a definite
number of protons in the nucleus, and that is what determines the identity of an
element. At this point in time, essentially only two elements, hydrogen and helium,
existed in the entire cosmos.
It is believed that there may also have been trace amounts of a third element,
known as lithium, among the first atoms formed after the big bang. Lithium is
identified by the presence of three protons along with four neutrons in the

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nuclei of its atoms. The three other atoms in the primeval universe are illustrated below.

–

+

++
–

–

Deuterium

Helium

D, He,
Li

+
++

–

–
–
Lithium

Chemistry makes use of a good amount of convenient symbolism. The elements
are represented by symbols that are usually the first one or two letters of their English,
Latin, or, in one case, German, names. Thus H stands for hydrogen, He for helium,
and Li for lithium. To represent the information about the content of an isotope’s
nucleus, we use isotopic notation. The number in the upper-left-hand corner is known
as the mass number, which is the total number of protons and neutrons. The number in the lower-left-hand corner is the atomic number, which is the number of protons in that specific element. The atomic number and the element’s symbol are
redundant because each element has a specific atomic number.
1
1H

2
1H

4
2He

7
3Li

The complete list of elements, along with their symbols, is shown inside the front
cover.
For a time after the formation of the original neutral atoms, the universe turned
completely dark. However, about 1 billion years after the big bang, some of the hydrogen and helium began to gather into individual clouds of gas. We are not sure why
this happened, but it did. However, this was a significant event because within the
regions of gas, an elementary attractive force of nature, gravity, began to have its way.
Gravity caused some of the individual clouds of gas to contract. Changes began to
occur in the clouds as the atoms moved closer together. As the clouds contracted,
the temperature within increased dramatically. Heat relates to the velocity of the
atoms. Heat causes expansion, which actually counteracts gravity. However, gravity
continued to predominate in the huge cloud, so further contraction and heating
occurred. Crushing pressures and high temperature began to build in the center of
the cloud.
In the dynamic, churning center of the cloud, the extremely hot hydrogen nuclei,
now stripped of their orbiting electrons, collided frequently with each other.
However, the fact that they had the same electrical charge meant that they actually

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A The Origin of Matter

did not touch but were repelled from each other
(like charges repel; unlike charges attract). At
some point, however, a phenomenal event happened in the evolution of our cosmos. Colliding
protons had enough velocity and energy to overcome the repulsive forces between the two nuclei,
thus merging or fusing together to form a single
nucleus. This process is called nuclear fusion. Some
of the mass of the two fusing nuclei converted
back into energy (i.e., Einstein’s law, E = mc2). It
was like starting a campfire. Once the fire starts,
it continues until the wood is consumed. The heat
generated from the fusion of the atoms within the
cloud was enough to counteract the pull of gravity, thus stabilizing the size of the cloud. The cloud
of gas began to glow and, at that time, a star was
born. The universe began to light up and appear
much like we see it today, except that the stars
were much closer together. (See Figure P-1.)
Fusion powers all of the stars, including our own
sun. (Nuclear fusion has only been understood
since the late 1930s.)

+

2

H

A Typical Star
Our sun is a typical star, generating
massive amounts of energy from
the fusion of hydrogen to form
helium.

FIGURE P-1

++

+

Deuterium
fusion

4
2He

2

H

The fusion of protons to form helium occurs in several steps that are somewhat
complex. The fusion of deuterium nuclei to form a helium nucleus, which was mentioned previously, however, occurs in one simple step and is illustrated below.
Many of the earliest stars were massive (many times larger than our own sun) and
used up their hydrogen fuel in the core of the star in the course of several million
years. Our own sun uses its fuel much more slowly, so even though it is over 4 billion years old, it will go on at this rate for many more billion years. So far, however,
in our early universe, the number of elements had not changed—just a minuscule
increase in the amount of helium as a result of fusion. It is what happened next in
the stars that led to the formation of heavier elements.
As the supply of hydrogen began to diminish in the core of the star, the fusion
of protons began to wane and no longer supplied enough energy to keep the size
of the star stable. So again gravity predominated and contraction recommenced.
With the increasing pressure came increasing temperatures. Eventually the heat was
high enough and the nuclei crushed together enough to cause the helium nuclei
(with a ⫹2 charge) to fuse to form the first elements heavier than lithium. (Stars in
this stage are known as red giants.) This process occurs at a higher temperature than
the proton fusion because of the higher nuclear charge to be overcome by the colliding nuclei. The fusion of helium in a series of steps eventually led to heavier elements such as carbon (mass number 12) and oxygen (mass number 16).
In time, the helium was exhausted and the contraction-heating process took over
once again. Another stage of elemental formation would begin, forming even heavier elements such as aluminum and silicon. Each stage of this cycle was a result of
higher temperatures and higher pressures within the star and led to the creation of
heavier and heavier elements.

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A Supernova
The bright star in the center of the
photo is actually an exploding star.
Heavy elements are being formed
in this explosion.

FIGURE P-2

13:41

This process of fusion, contraction, higher temperature, and pressure followed
by more fusion only goes so far. The fusing of lighter elements to form elements
up to the mass of iron (26 protons and 30 neutrons) releases energy. Formation of
any element heavier than that does not release energy, so the fusion process stops
abruptly at that point. It is as if iron forms the ashes of the nuclear fire. When a
fire is reduced to ashes, the fire goes out. However, in the stepwise evolution of a
star, the first 26 or 27 elements were formed. Yet we have many other familiar elements that are much heavier than iron, such as gold and uranium. What is their
origin if not from the interior of a star? The spectacular results of what happens
when a huge star exhausted the fuel in its now mainly iron core is what produced
the heaviest of the elements.
Other messages from space have given us insight into creation of the heavier
elements. In 1987, an extremely bright star was found in the heavens from an
observatory in Chile. Only an ordinary dim star had been there the night before.
The astronomers became immediately aware that they were witnessing the explosion of a star known as a supernova. (See Figure P-2.) This phenomenon had been
previously known but was now being witnessed. Apparently, the star had suddenly
exhausted the supply of fuel for fusion in its core. Without an energy source to counteract gravity, the star again collapsed. This time the collapse did not stop, causing
densities, temperatures, and pressures to go to the extreme. But in a matter of seconds, the collapsing star rebounded in a cataclysmic explosion, propelling the outer
layers of the star into outer space along with a huge flux of neutrons (formed when
electrons are forced into a proton). Elements such as iron dramatically increased
their mass by adding neutrons during and after the explosion. The more massive
atoms formed at this time eventually underwent nuclear changes that produced
elements with higher and higher atomic numbers such as silver, lead, and gold.
To summarize the sequence of events: (1) A star explodes, sending elements
formed by fusion in the star rushing into space along with huge numbers of neutrons. (2) The atoms of the original elements absorb neutrons, increasing the atomic
mass. (3) The heavier elements undergo nuclear changes, leading to elements with
high atomic numbers. (The heaviest naturally occurring element found on Earth is
uranium-238.)
For billions of years, giant clouds of hydrogen have formed and contracted into new
generations of stars. Each generation had more and more of the heavier elements.
(In fact, astronomers have identified stars with few heavy elements, indicating that they
date to near the beginning of time.) Blasted by the force of supernovae, this matter
drifted through cold, dark space. About 4.6 billion years ago, a cloud of hydrogen containing the dust and debris from previous stars condensed to form our sun, with
enough star dust left over to form eight (or more) planets and many other smaller
solid bodies trapped in perpetual orbits around this star. On the third planet out from
this ordinary star, atoms of some of the lighter elements assembled into living systems.
It sounds somewhat melodramatic, but we are indeed children of the stars.

The Origin of the Chemical Bond
So far, we have encroached into the disciplines of astronomy and physics. Chemistry
came into being when atoms began to bond to each other. Let’s return to the early
universe, about the time clouds of neutral hydrogen atoms (nuclei and orbiting electrons) were beginning to form. At first, the new hydrogen atoms were very hot. When
two hydrogen atoms collided, they bounced away from each other like recoiling billiard balls. When they cooled sufficiently, however, another important event
occurred. The two neutral atoms stuck together. The electrons of the two atoms acted
like glue, holding the two atoms together.
The sharing of electrons between two atoms is known as a covalent chemical bond
or simply a chemical bond. If it weren’t for the fact that two atoms are more stable
bonded together than existing apart, the universe would still be just a collection of

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B The Mystery of Fire

9

gaseous atoms—no continents, no oceans, no moons, no us. The two hydrogen
atoms joined together to form a molecule. A molecule is composed of two or more
atoms chemically bonded together. This is illustrated by the formula H2, where the
subscript 2 indicates the number of hydrogen atoms present in the molecule. The
hydrogen molecule is the primeval molecule. It is the way hydrogen exists except at
a very high temperature that forces the atoms apart.
As the universe became enriched in other elements, chemical bonds formed
between other elements. In the farthest reaches of space, we have detected many
other molecules that have formulas such as H2CO (formaldehyde whose molecules
contain two hydrogens, one carbon, and one oxygen), HCN (hydrogen cyanide),
H2O (water), NH3(ammonia), and even more complex species. Like the primeval
molecule, H2, the atoms in these molecules are held together by chemical bonds.
(There is another type of bond that held lithium and hydrogen together in the early
universe called an ionic bond. It will be described in Chapter 2.) On Earth, we find
molecules ranging from the simplest two-atom hydrogen molecule to those of DNA,
which contain millions of atoms.
In Chapter 1, we begin our journey into chemistry by emphasizing mathematical
skills related to chemistry. In Chapter 2, we then examine the types of matter found
in our beautiful planet with its air, oceans, and solid earth. We will then develop and
master the topic one step at a time while continuing to emphasize the “big picture.”

B

THE MYSTERY OF FIRE

At about the same time (give or take 100,000 years) that humans may have started
to wonder about the meaning of the night sky, they did something more immediately practical with their new ability to reason. It would change their destiny and
order in the realm of all other animals. They tamed and put to use one of the most
powerful forces in nature: fire.
It is difficult to imagine how our ancient ancestors could have managed without
fire. Humans do not have sharp night vision like the raccoon, but fire brought light
to the long, dark night. We have no protective fur like the deer, but fire lessened the
chill of winter. We do not have sharp teeth or powerful jaws like the lion, but fire rendered meat tender. Humans are not as strong or as powerful as the other large animals, but fire repels even the most ferocious of beasts. It seems reasonable to suggest
that the taming of fire was one of the most monumental events in the history of the
human race. The use of fire made our species dominant over all others.
Let’s fast-forward in time to near the end of the Stone Age, about
10,000 years ago, when fire became the agent that launched us into the
world of chemistry. In the Stone Age, weapons and utensils were fashioned from rocks and a few chunks of copper metal (an element) that
were found in nature. Copper was superior to stone because pounding
could easily shape it into fine points and sharp blades. Unfortunately,
native copper was quite rare. But about 7000 years ago this changed.
Anthropologists speculate that some resident of ancient Persia found copper metal in the ashes remaining from a hot charcoal fire. The free copper had not been there before, so it must have come from a green stone
called malachite (see Figure P-3), which probably lined the fire pit.
Imagine the commotion that this discovery must have caused. Hot coals
could transform a particular stone into a valuable metal. Fire was the key
that launched the human population into the age of metals. The recovery of metals from their ores is now a branch of chemical science called F I G U R E P - 3 Malachite Malachite is a
metallurgy. The ancient Persians must have considered this discovery a dra- copper ore. When heated with charcoal, it
matic example of the magic of fire.
forms metallic copper.

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Other civilizations used chemistry in various ways. About 3000 B.C., the Egyptians
learned how to dye cloth and embalm their dead through the use of certain chemicals found in nature. They were very good at what they did. In fact, we can still determine from ancient mummies the cause of death and even diseases the person may
have had. The Egyptians were good chemists, but they had no idea why any of these
procedures worked. Every chemical process they used was discovered by accident.
Around 400 B.C., while some Greeks were speculating about their various gods,
philosophers were trying to understand and describe nature. These great thinkers
argued about why things occurred in the world around them, but they were not
inclined (or able) to check out their ideas by experimentation or to put them to
practical use. At the time, however, people believed that there were four basic elements of nature—earth, air, water, and fire. Everything else was simply a specific combination of these basic elements. Of the original four elements, fire was obviously
the most mysterious. It was the transforming element; that is, it had the capacity to
change one substance into another (e.g., certain rocks into metals). We now call such
transformations “chemistry.” Fire itself consists of the hot, glowing gases associated
with certain chemical changes. If fire is a result of an ongoing chemical transformation, then it is reasonable to suggest that chemistry and many significant advances
in the human race are very much related.
The early centuries of the Middle Ages (A.D. 500–1600) in Europe are sometimes
referred to as the Dark Ages because of the lack of art and literature and the decline
of central governments. The civilizations that Egypt, Greece, and Rome had previously built began to decline. Chemistry, however, began to grow during this period, especially in the area of experimentation. Chemistry was then considered a combination
of magic and art rather than a science. Many of those who practiced chemistry in
Europe were known as alchemists. Some of these alchemists were simply con artists who
tried to convince greedy kings that they could transform cheaper metals such as lead
and zinc into gold. Gold was thought to be the perfect metal. Such a task was impossible, of course, so many of these alchemists met a drastic fate for their lack of success.
However, all was not lost. Many important laboratory procedures such as distillation
and crystallization were developed. Alchemists also discovered or prepared many previously unknown chemicals, which we now know as elements and compounds.
Modern chemistry has its foundation in the late 1700s when the use of the analytical balance became widespread. Chemistry then became a quantitative science
in which theories had to be correlated with the results of direct laboratory experimentation. From these experiments and observations came the modern atomic theory, first proposed by John Dalton around 1803. This theory, in a slightly modified
form, is still the basis of our understanding of nature today. Dalton’s theory gave
chemistry the solid base from which it could serve humanity on an impressive scale.
Actually, most of our understanding of chemistry has evolved in the past 100 years.
In a way, this makes chemistry a very young science. However, if we mark the beginning of chemistry with the use of fire, it is also the oldest science.
From the ancient Persians five millennia ago, to the Egyptians, to the alchemists
of the Middle Ages, various cultures have stumbled on assorted chemical procedures.
In many cases, these were used to improve the quality of life. With the exception of
the Greek philosophers, there was little attention given to why a certain process
worked. The “why” is very important. In fact, the tremendous explosion of scientific
knowledge and applications in the past 200 years can be attributed to how science
is now approached. This is called the scientific method, which we will discuss next.

C

THE SCIENTIFIC METHOD

In ancient times, scientific advances were discovered by accident. This still occurs
to some extent, but we have made great strides in how we approach science such
that most modern advances occur by design. Consider the case of tamoxifen, a chem-

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11

ical that is saving the lives of thousands of women in the United States alone. It is
thought to dramatically reduce the chances of recurrence of breast cancer, a frightening disease that now strikes about one in nine American women. In 1998, the Food
and Drug Administration (FDA) approved use of the drug to prevent breast cancer
in women with a high risk of the disease. But the drug also has undesirable side
effects. It may increase the chances of other types of cancer in some women and, in
a few cases, causes depression, irritability, and short-term memory loss. However,
knowing how tamoxifen works led scientists to produce an improved drug. In
December 2004, the drug anastrazole was found to be at least as effective as tamoxifen but with fewer side effects. It is now under large-scale testing. How tamoxifen
was discovered in the first place and what led to the potentially improved drug are
examples of how the scientific method improves our lives. But the first step in the scientific method is a long way from producing a useful drug. It simply involves making observations and gathering data. As an example, imagine that we are the first to
make a simple observation about nature—“the sun rises in the east and sets in the
west.” This never seems to vary and, as far as we can tell from history, it has always
been so. In other words, our scientific observation is strictly reproducible. So now we
ask “Why?” We are ready for a hypothesis. A hypothesis is a tentative explanation of
observations. The first plausible hypothesis to explain our observations was advanced
by Claudius Ptolemy, a Greek philosopher, in A.D. 150. He suggested that the sun,
as well as the rest of the universe, revolves around the Earth from east to west. That
made sense. It certainly explained the observation. In fact, this concept became an
article of religious faith in much of the Western world. However, Ptolemy’s hypothesis did not explain other observations known at the time, which included the movement of the planets across the sky and the phases of the moon.
Sometimes new or contradictory evidence means a hypothesis, just like a brokendown old car, must either receive a major overhaul or be discarded entirely. In 1543,
a new hypothesis was proposed. Nicolaus Copernicus explained all of the observations
about the sun, moon, and planets by suggesting that Earth and the other planets orbit
around the sun instead of vice versa. Even though this hypothesis explained the mysteries of the heavenly bodies, it was considered extremely radical and even heretical
at the time. (It was believed that God made
Earth the center of the universe.) In 1609,
a Venetian scientist by the name of Galileo
Galilei built a telescope to view ships still
far out at sea. When he turned the telescope up to the sky, he eventually produced almost unquestionable proof that
Copernicus was correct. Galileo is sometimes credited with the beginning of the
modern scientific method because he provided direct experimental data in support
of a concept. The hypothesis had withstood the challenge of experiments and
thus could be considered a theory. A theory is a well established hypothesis. A theory
should predict the results of future experiments or observations.
The next part of this story comes in
1684, when an English scientist named Sir
Isaac Newton stated a law that governs the
motion of planets around the sun. A law
is a concise scientific statement of fact to which
no exceptions are known. Newton’s law of
universal gravitation states that planets
are held by gravity in stationary orbits F I G U R E P - 4 The Solar System The Copernican theory became the basis of
a natural law of the universe.
around the sun. (See Figure P-4.)

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In summary, these were the steps that led to a law of nature:
1. Reproducible observations (the sun rises in the east)
2. A hypothesis advanced by Ptolemy and then a better one by Copernicus
3. Experimental data gathered by Galileo in support of the Copernican
hypothesis and eventual acceptance of the hypothesis as a theory
4. The statement by Newton of a universal law based on the theory
Variations on the scientific method serve us well today as we pursue an urgent
search for cures of diseases. An example follows.

The Scientific Method in Action

The Pacific
Yew The bark of this tree is a
source of an anticancer drug
known as Taxol.

FIGURE P-5

The healing power of plants and plant extracts has been known for thousands of
years. For example, ancient Sumerians and Egyptians used willow leaves to relieve
the pain of arthritis. We now know that extracts of the common willow contain a
drug very closely related to aspirin. This is the observation that starts us on our journey to new drugs. An obvious hypothesis comes from this observation, namely, that
there are many other useful drugs among the plants and soils of the world. We
should be able to find them. There are several recent discoveries that support this
hypothesis. For example, the rosy periwinkle is a common tropical plant not too
different from thousands of other tropical plants except that this one saves lives.
The innocent-looking plant contains a powerful chemical called vincristine, which
can cure childhood leukemia. Another relatively new drug called Taxol has been
extracted from the Pacific yew tree. (See Figure P-5.) Taxol is effective in treating
ovarian cancer and possibly breast cancer. Others include cyclosporine, isolated
from a fungus in 1957, which made organ transplants possible, and digoxin, isolated
from the foxglove plant, used for treatment of heart failure. In fact, many of the
best-selling medicines in the United States originated from plants and other natural sources. Besides those mentioned, other drugs treat conditions such as high
blood pressure, cancer, glaucoma, and malaria. The search for effective drugs from
natural sources and newly synthesized compounds is very active today. These chemicals are screened for potential anticancer, antiarthritic, and anti-AIDS activity. Since
the greatest variety of plants, molds, and fungi are found in tropical forests, these
species are receiving the most attention. The introduction of a new medicine from
a plant involves the following steps.
1. Collection of materials. “Chemical prospectors” scour the backwoods of the
United States and the tropical forests such as those in Costa Rica, collecting
and labeling samples of leaves, barks, and roots. Soil samples containing
fungi and molds are also collected and carefully labeled.
2. Testing of activity. Scientists at several large chemical and pharmaceutical
companies make extracts of the sample in the laboratory. These extracts are
run through a series of chemical tests to determine whether there is any
antidisease activity among the chemicals in the extract. New methods such
as high-throughput-screening (HTS) allow certain chemical companies to
screen thousands of chemicals in a single day. If there is antidisease activity,
it is considered a “hit” and the extract is taken to the next step.
3. Isolation and identification of the active ingredient. The next painstaking task is
to separate the one chemical that has the desired activity from among the
“soup” of chemicals present. Once that’s done, the particular structure of
the active chemical must be determined. A hypothesis is then advanced
about what part of the structure is important and how the chemical works.
The hypothesis is tested by attempting to make other more effective drugs
(or ones with fewer side effects) based on the chemical’s structure. It is then
determined whether the chemical or a modified version of it is worth
further testing.

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4. Testing on animals. If the chemical is considered promising, it is now ready to
be tested on animals. This is usually done in government and university labs
under strictly controlled conditions. Scientists study toxicity, side effects, and
the chemical’s activity against the particular disease for which it is being tested.
If, after careful study, the chemical is considered both effective and safe, it is
ready for the next step.
5. Testing on humans. The final step is the careful testing on humans in a series
of clinical trials carefully monitored by agencies such as the FDA.
Effectiveness, dosage, and long-term side effects are carefully recorded and
evaluated. All told, it currently takes from 10 to 15 years for a new drug to
make it all the way from research and development to market.
If a chemical with the desired activity is randomly discovered, only about one in
1000 may actually find its way into general use. Still, the process works. Many chemicals active against cancer and even AIDS (one has been isolated from the mulberry
tree) are now in the pipeline for testing. There is some urgency in all of this. Not
only are we anxious to cure specific diseases, but the tropical forests that contain the
most diverse plants are disappearing at an alarming rate. In any case, nature is certainly our most important chemical laboratory.
At this time, most of the drugs that originated from synthetic or natural sources
are considered cases of “serendipity.” We are now moving more toward the concept
of “rational drug design.” Here, the goal is to identify the active sites on the molecules of diseases such as viruses, tumors, or bacteria. The next step is to deliberately
synthesize a drug that attaches to that active site and either destroys or otherwise alters
the disease molecules. This sounds easy, but it is not. It requires that we know more
about the structure and geometry of these disease molecules. We can then advance
hypotheses as to how designed molecules would interact. This is the direction in which
pharmaceutical chemists are heading, however.
The scientific method has produced cures for many diseases that have plagued
the human race for thousands of years. Eventually, it may lead to a cure or vaccine
for AIDS. Our modern scientific methods also produce materials for space travel,
all sorts of plastics and synthetic fibers, microchips for computers, and processes for
genetic engineering. Just a century ago, everything was made out of stone, wood,
glass, metal, or natural fibers (wool, cotton, and silk). Our modern society could
hardly function on those materials alone.

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C H A P T E R

1
Measurements
in Chemistry

T

o help fly and land this large aircraft safely, a huge number of dials and
gauges are displayed in the cockpit. Ground speed, temperature, cabin
pressure, altitude, and location are just a few of the measurements
that must be continually available to the experienced
pilot. To a large degree, the study of nature in general
and chemistry in particular is based on observations. But there is much more. We require
reproducible, quantitative measurements.
How measurements in chemistry are
expressed and manipulated are the
subjects of this chapter.

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PA R T A
THE NUMBERS USED IN
CHEMISTRY

1-1 The Numerical Value of a
Measurement

1-2

Significant Figures and
Mathematical Operations
MAKING IT REAL

Ted Williams and Significant Figures

1-3

Expressing Large and Small
Numbers: Scientific Notation

PA R T B
THE MEASUREMENTS USED
IN CHEMISTRY

1-4

Measurement of Mass, Length,
and Volume
MAKING IT REAL

Worlds from the Small to the Distant—Picometers
to Terameters

1-5

Conversion of Units by the
Factor-Label Method

1-6 Measurement of Temperature

S E T T I N G T H E S TA G E Before we were even
aware of our own existence, we were being subjected to
measurements. Within the first few minutes of our lives, we
were placed on a scale and against a tape measure to provide our first weight and length. Years later, as adults, we
now find ourselves immersed in a complex world that measures just about anything that lends itself to being measured.
For example, it is hard to conceive of being able to operate a
modern automobile without a speedometer as well as temperature and fuel gauges. The more expensive the car, the
more measurements that are reported by more gauges and
dials. If one pays enough, a global positioning satellite will
even measure where you are and tell you how to get where
you are going. Pilots, carpenters, artists, teachers, and most
other professionals and skilled workers are usually involved
in some type of measurement. So chemists are hardly
unique in their reliance on measurements.
Chemistry is a science that requires us to deal with all the
stuff that we see around us. This stuff, which we call matter, is
subject to the laws of nature. Many of these laws originate
from reproducible quantitative measurements. Measurements
naturally contain numbers. The quality, meaning, magnitude,
and manipulation of the numbers we use in chemistry form
the beginning point of our study. Our journey into the actual
concepts of chemistry begins in the next chapter.
If you are not in at least fair physical shape, playing a
strenuous sport is not fun—it is exhausting and probably
frustrating. Likewise, doing chemistry can be frustrating if
you are not in at least fair mathematical shape. It is likely
that most students need at least a bit of a mathematical
workout to get into shape. For some, a simple review will do,
and that is provided in the text. At appropriate points in this
chapter, you will also be referred to a specific appendix in
the back of the book for more extensive review. Review
appendixes include basic arithmetical operations (Appendix A),
basic algebra operations (Appendix B), and scientific
notation (Appendix C). Also, Appendix E can aid you in the
use of calculators for the mathematical operations found in
the text. If you are worried about the math, you are not
alone. Just remember that this text was written with your
concerns in mind.
Measurements consist of two parts—a number and a
specific unit. We will discuss these two parts separately.
In Part A, we will address questions about the numerical
quantity of a measurement. The units that are used in the
measurement are then discussed in Part B.

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PA R T A
THE NUMBERS USED IN
CHEMISTRY

OBJ ECTIVES

SETTI NG A GOAL
■

You will learn how to apply and manipulate
measurements to produce scientifically meaningful
outcomes.

1-1 (a) Describe the difference between accuracy and
precision. (b) Determine the number of significant figures
in a measurement.
1-2 Perform arithmetic operations, rounding the answer to
the appropriate number of significant figures.
1-3 (a) Write very large or small measurements in scientific
notation. (b) Perform arithmetic operations involving
scientific notation.

䉴 OBJECTIVES

FOR

SECTION 1-1

(a) Describe the difference between
accuracy and precision. (b) Determine
the number of significant figures in a
measurement.

1-1

T H E N U M E R I C A L VA L U E O F A M E A S U R E M E N T
A H E A D ! Much of science is based on numbers. How reliable
are the numbers and what do they really tell us? In this section, we will
evaluate the quality and reliability of the numbers that are part of a
measurement. ■
LOOKING

1-1.1 The Qualities of a Number
We will start this chapter with a formal definition of a measurement. A measurement
determines the quantity, dimensions, or extent of something, usually in comparison to a specific unit. A unit is a definite quantity adopted as a standard of measurement. Thus, a measurement (e.g., 1.23 meters) consists of two parts: a numerical quantity (1.23)
followed by a specific unit (meters).
First let’s consider the numerical value of a measurement. Assume that the evening
news informs us that 12,000 people gathered for a concert. Did they mean exactly 12,000?
Not really—actually, it was just an estimate. Two other experts may have estimated the
same crowd at 13,000 and 11,000, respectively. This means that the original estimate had
an uncertainty of ⫾1000. Thus only the 1 and 2 are considered significant in this estimate. The three zeros simply tell us the magnitude of the number. In a measurement,
a significant figure is a digit that is either reliably known or closely estimated. In the number
12,000, we can assume the 1 is reliable and reproducible from any number of estimates,
but the 2 is estimated. The zeros are not significant since they actually have no specific
numerical meaning. Thus, in our example, there are two significant figures: the 1 and
the 2. The number of significant figures or digits in a measurement is simply the number of measured digits and refers to the precision of the measurement. Precision relates
to the degree of reproducibility or uncertainty of the measurement. Indeed, all measured values
have an uncertainty that is expressed in the last significant figure to the right.
Now assume the same crowd at the concert is seated in the bleachers of a stadium instead of milling about. In this case, a more precise estimate is possible since
the exact capacity of the stadium is known. The crowd can now be estimated at
between 12,400 and 12,600, or an average of 12,500. This is a measurement with
three significant figures. The 1 and 2 are now reliable, but the third significant figure, the 5, is estimated. The extra significant figure means that the uncertainty is
now reduced to ⫾100. The more significant figures in a measurement, the more
precise it is. If the crowd went through a turnstile before entering the stadium, an
even more precise number could be given. Notice that the significant figure farthest
to the right in a measurement is estimated.

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17

Participants in the sport of riflery
(target shooting) are judged on two
points: how close the bullet holes are to
each other (the pattern) and how close
the pattern is to the center of the target known as the bull’s-eye. The precision of the contestant’s shooting is
measured by the tightness of the pattern. How close the pattern is to the
High accuracy
Low accuracy
Low accuracy
bull’s-eye is a measure of the shooter’s
High precision
Low precision
High precision
accuracy. Accuracy in a measurement
refers to how close the measurement is to the true value. Usually, the more precise the meas- F I G U R E 1 - 1 Precision and
urement, the more accurate it is—but not always. In our example, if a certain com- Accuracy High precision does not
petitor has a faulty sight on the rifle, the shots may be close together (precise) but necessarily mean high accuracy.
off center (inaccurate). (See Figure 1-1.)
Accuracy in measurements depends on how carefully the instrument of measurement has been calibrated (compared to a reliable standard). For example, what if
we attempted to measure length with a plastic ruler that became warped after being
left in the hot sun? We obviously would not obtain accurate readings. We would need
to recalibrate the ruler by comparing its length divisions to a reliable standard.

1-1.2 Zero as a Significant Figure
It would be easy to determine the number of significant figures in a measurement
if it were not for the number zero. Unfortunately, zero serves two functions: as a reliable or estimated digit, or simply as a marker to locate the decimal point (such as
the three zeros in the estimated crowd of 12,000 people). Since the zeros look alike
in both cases, it is important for us to know whether a zero is significant or is there
simply to locate the decimal point. The following rules can be used to tell us about
zero. Digits that are underlined are significant.
1. When a zero is between other nonzero digits, it is significant. 709 has three
significant figures. Zeros between two nonzero digits are always siginificant.
2. Zeros to the right of a nonzero digit and to the right of the decimal point
are significant. 8.0 has two significant figures, just as 7.9 and 8.1 do. 7.900
has four significant figures. There is no value difference between 8.0 and
just 8. Why, then, is the zero added? Quite simply because it is an estimated
digit and should be included.
3. Zeros to the left of the first nonzero digit are not significant. 0.0078 and
0.45 both have two significant figures; 0.04060 has four significant figures.
In this case, the zeros to the left of the digit are simply showing the decimal
place and are not measured or estimated digits. Therefore, they are not
significant.
4. Zeros to the left of an implied decimal point may or may not be significant.
In most cases, they are not. The crowd of 12,000 has two significant figures,
as does 6600. Just as in the above case, the zeros to the left are decimal place
holders, not actual measurements. What if the zero in a number such as 890
is actually an estimated digit and thus significant? This is a tough question.
Some texts use a line over the zero to indicate that it is significant (e.g., 890),
and others simply place a decimal point after the zero (e.g., 890.). As we will
see in Section 1-3, there is a solution to this dilemma.
In most problems used in this text, measurements are expressed to three significant figures. Therefore, in calculations where other numbers have three
significant figures, we will assume that numbers such as 890 also have three
significant figures.

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1-1

Evaluating Zero as a Significant Figure
How many significant figures are in the following measurements? What is the uncertainty in each of the measurements?
(a) 1508 cm
(b) 300.0 ft
(c) 20.003 lb
(d) 0.00705 gal
PROCEDURE

To determine the number of significant figures, refer to the rules regarding zero that were listed. Since these numbers all involve
measurements, the last significant figure to the right is estimated. This digit indicates the uncertainty of the measurement.
SOLUTION

(a) The zero between the two nonzero digits is significant. There are four significant figures. The uncertainty is in the last
digit to the right, so is 1 cm.
(b) The zero to the right of the decimal point is significant because it is displayed. Therefore, the other two zeros are
also significant because they lie between significant figures. There are four significant figures. The uncertainty is in
the last zero to the right, so is 0.1 ft.
(c) All the zeros are between nonzero digits, so they are all significant. There are five significant figures. The uncertainty is
in the last digit to the right, so is 0.001 lb.
(d) The three zeros to the left of the first nonzero digit are not significant. The zero between the two nonzero digits is significant. Thus, there are three significant figures. The uncertainty is in the last digit to the right, so is 0.00001 gal.
A N A LY S I S

Be certain not to confuse the concept of “significant” with the idea of “important.” In the measurement 0.00705 gallons, the
two zeros immediately following the decimal are most assuredly important to the value of the number. Without them, the value
changes. But significant refers to a measurement. In that particular example, the measurement doesn’t start until you reach
the 7. Therefore, the zeros, while necessary, are not significant.
SYNTHESIS

The more significant figures there are in a measurement, the more precise it is. The closest the Earth and sun ever get to one
another is 147,098,000 km. The uncertainty is 1000 km. The distance between New York and Philadelphia is 127 km, known
with a precision to the nearest kilometer. One would assume that knowing a distance to a kilometer would be more precise
than to the nearest 1000 kilometers. And yet, because the measured distance to the sun has six significant figures in it, it is
1000 times more precise than the measured driving distance between the two cities. Which of the four measurements in the
problem is the most precise? 20.003 lb. It has more significant figures than any of the others.

ASSESSING THE
OBJECTIVES FOR
SECTION 1-1

1 - 1 ( a ) K N O W L E D G E : Students in different class sections
attempted to measure the mass of a 100.0-gram object. Describe the series of
measurements as precise, precise and accurate, or neither.
(a) class 1: 94.2 g, 93.8 g, 94.4 g, 94.0 g
(b) class 2: 94.3 g, 89.7 g, 102.4 g, 97.8 g
(c) class 3: 100.2 g, 100.0 g, 99.8 g, 99.8 g

EXERCISE

You are determining the boiling point of an
unknown compound in degrees Celsius. On three successive attempts, you record
145.5°C, 145.8°C, and 146.0°C. Are your data precise? Are your data accurate? Is
there extra information needed to make either of those claims?

EXERCISE 1-1(b) SYNTHESIS:

E X E R C I S E 1 - 1 ( c ) A N A LY S I S :

in the following measurements:
(a) 45.00 oz
(b) 0.045 lbs

Determine the number of significant figures

(c) 45000 s

(d) 0.450 in.

1 - 1 ( d ) S Y N T H E S I S : Two models of analytical balance are
available for purchase. The first measures to  0.01 gram. The second is more

EXERCISE

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19

expensive, but measures to  0.0001 gram. How many significant figures will each
give when measuring the mass of an object of about 25 grams? What considerations
would you include in determining which model to buy?
(Throughout the text, answers to all Assessing the Objectives exercises can be found at the
end of each chapter.)
For additional practice, work chapter problems 1-1, 1-3, 1-4, and 1-6.

1-2

SIGNIFICANT FIGURES AND
M AT H E M AT I C A L O P E R AT I O N S

䉴 OBJECTIVE

FOR
SECTION 1-2

Perform arithmetic operations, rounding the answer to the appropriate
number of significant figures.

L O O K I N G A H E A D ! In science we are continually adding, multiplying, or
performing other calculations involving various measurements. Often these
measurements involve different degrees of precision, so we need to understand
how to report the results of such calculations. ■

For under $10 we can own a small calculator that can almost instantly carry out any
calculation we may encounter in general chemistry. This device is truly phenomenal, especially to older scientists, who years ago had to carry out these calculations
using a slide rule. (This was a simple but effective device that performed many of
the calculations found on an electronic calculator but not as quickly or as precisely.)
The use of the calculator does have one serious drawback, however. It does not necessarily report the answer to a calculation involving measurements to the proper precision or number of significant figures. For example, 7.8 divided by 2.3 reads
3.3913043 on a standard eight-digit display. However, if the numbers represented
measurements known to only two significant figures (e.g., 7.8 lb and 2.3 qt), then
the calculator has created the illusion that these numbers were known to a much
greater precision. Since the calculator does not express the answer to the calculation with the proper precision, we must know how to prune the answer so that the
reported value is honest and appropriate.

1-2.1 Rules for Addition and Subtraction and Rounding Off
There are two sets of rules for properly expressing the result of a mathematical
operation: one applies to addition and subtraction and the other applies to multiplication and division. We will discuss the rule for addition and subtraction first.
When numbers are added or subtracted, the answer is expressed to the same number of decimal places as the measurement with the fewest decimal places. Or, in other
words, the summation must have the same degree of uncertainty as the measurement
with the most uncertainty (e.g., 100 has more uncertainty than 10 and 0.1 more
than 0.01). This is illustrated by the following summation.
10.6871
1.42
12.1071

=

(four decimal places, or uncertainty of ; 0.0001)
(two decimal places, or uncertainty of ; 0.01)
12.11(two decimal places, or uncertainty of ; 0.01)

Notice that 1.42 and the answer have the same uncertainty. Therefore, expressing the
71 in the summation has no meaning and cannot be included except for rounding-off
purposes. The calculator does not give the answer to the proper number of decimal
places.
The rules for rounding off a number are as follows (the examples are all to be
rounded off to three significant figures).
1. If the digit to be dropped is less than 5, simply drop that digit (e.g., 12.44 is
rounded down to 12.4).

The calculator does not give the
answer to the proper number of decimal places or significant figures.

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2. If the digit to be dropped is 5 or greater, increase the precedin