Preview Introductory chemistry an atoms first approach by Julia Burdge Michelle Driessen (2017)

Preview Introductory chemistry an atoms first approach by Julia Burdge Michelle Driessen (2017) Preview Introductory chemistry an atoms first approach by Julia Burdge Michelle Driessen (2017) Preview Introductory chemistry an atoms first approach by Julia Burdge Michelle Driessen (2017) Preview Introductory chemistry an atoms first approach by Julia Burdge Michelle Driessen (2017) Preview Introductory chemistry an atoms first approach by Julia Burdge Michelle Driessen (2017)

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Library of Congress Cataloging-in-Publication Data

Names: Burdge, Julia | Driessen, Michelle.

Title: Introductory chemistry : an atoms first approach / Julia Burdge,

Michelle Driessen.

Description: First edition | New York, NY : McGraw-Hill, 2015.

Identifiers: LCCN 2015040623| ISBN 9780073402703 (alk paper) | ISBN

0073402702 (alk paper)

Subjects: LCSH: Chemistry—Textbooks.

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To the people who will always matter the most: Katie, Beau, and Sam.

—Julia Burdge

To my family, the center of my universe and happiness, with special thanks to my husband for his support and making me the person I am today.

—Michelle Driessen

And to Robin Reed, for her timely and hilarious memes—and for her eternal good humor.

—Julia Burdge and Michelle Driessen

About the Authors

Julia Burdge holds a Ph.D (1994) from The University of Idaho in Moscow, Idaho; and a Master’s Degree from The University of South Florida

Her research interests have included synthesis and characterization of cisplatin analogues, and development of new analytical techniques and instrumentation for measuring ultra-trace levels of atmospheric sulfur compounds

She currently holds an adjunct faculty position at The College of Western Idaho in Nampa, Idaho, where she teaches general chemistry using an atoms first approach; but spent the lion’s share of her academic career at The University of Akron in Akron, Ohio, as director of the Introductory Chemistry program In addition to directing the general chemistry program and supervising the teaching activities of graduate students, Julia established a future-faculty development program and served as a mentor for graduate students and post doctoral associates

Julia relocated back to the Northwest to be near family In her free time, she enjoys precious time with her three children, and with Erik Nelson, her partner and best friend

Michelle Driessen earned a Ph.D in 1997 from the University of Iowa in Iowa City, Iowa Her research and dissertation focused on the thermal and photochemical reactions of small molecules at the surfaces of metal nanoparticles and high surface area oxides

Following graduation, she held a tenure-track teaching and research position

at Southwest Missouri State University for several years A family move took her back to her home state of Minnesota where she held positions as adjunct faculty at both St Cloud State University and the University of Minnesota It was during these adjunct appointments that she became very interested in chemical education Over the past several years she has transitioned the general chemistry laboratories at the University of Minnesota from verification

to problem-based, and has developed both online and hybrid sections of general chemistry lecture courses She is currently the Director of General Chemistry at the University of Minnesota where she runs the general chemistry laboratories, trains and supervises teaching assistants, and continues to experiment with active learning methods in her classroom

Michelle and her husband love the outdoors and their rural roots They take every opportunity to visit their family, farm, and horses in rural Minnesota

Brief Contents

1 Atoms and Elements 2

2 Electrons and the Periodic Table 30

3 Compounds and Chemical Bonds 74

4 How Chemists Use Numbers 122

5 The Mole and Chemical Formulas 162

6 Molecular Shape 192

7 Solids, Liquids, and Phase Changes 234

8 Gases 268

9 Physical Properties of Solutions 308

10 Chemical Reactions and Chemical Equations 344

11 Using Balanced Chemical Equations 382

12 Acids and Bases 416

1 ATOMS AND ELEMENTS 2

1.1 The Study of Chemistry 3

• Why Learn Chemistry? 3

• The Scientific Method 3

1.2 Atoms First 5

1.3 Subatomic Particles and the

Nuclear Model of the Atom 6 1.4 Elements and the Periodic Table 10

■ Elements in the Human Body 11

1.5 Organization of the Periodic Table 14

■ Elements in Earth’s Crust 15

1.6 Isotopes 16

■ Mass Spectrometry 17

■ Iron-Fortified Cereal 19

1.7 Atomic Mass 19

2.1 The Nature of Light 31

2.7 Ions: The Loss and Gain of Electrons 61

• Electron Configuration of Ions 61

• Lewis Dot Symbols of Ions 63

Contents

vi

3 COMPOUNDS AND CHEMICAL BONDS 74

3.1 Matter: Classification and Properties 75

• States of Matter 75 • Mixtures 76

• Naming Atomic Cations 86

• Naming Atomic Anions 87

• Naming Binary Ionic Compounds 87

3.4 Covalent Bonding and Molecules 89

• Covalent Bonding 90 • Molecules 90

• Molecular Formulas 93

■ Fixed Nitrogen in Fertilizers 96

3.5 Naming Binary Molecular Compounds 97 3.6 Covalent Bonding in Ionic Species: Polyatomic Ions 99

■ Product Labels 100

■ Product Labels 101

3.7 Acids 105 3.8 Substances in Review 107

Properties of Atoms 108

• Distinguishing Elements and Compounds 110

• Determining Whether a Compound Is Ionic or Molecular 111

• Naming Compounds 111

vii

4 HOW CHEMISTS USE NUMBERS 122

4.1 Units of Measurement 123

• Base Units 123 • Mass, Length, and Time 124

• Metric Multipliers 124 • Temperature 126

■ The Fahrenheit Temperature Scale 127

4.2 Scientific Notation 130

• Very Large Numbers 131 • Very Small Numbers 132 • Using the Scientific Notation Function on Your Calculator 133

4.5 Success in Introductory Chemistry Class 152

5.1 Counting Atoms by Weighing 163

• The Mole (The “Chemist’s Dozen”) 163

• Molar Mass 165 • Interconverting Mass, Moles, and Numbers of Atoms 167

5.2 Counting Molecules by Weighing 169

• Calculating the Molar Mass of a Compound 169 • Interconverting Mass, Moles, and Numbers of Molecules (or Formula Units) 170 • Combining Multiple Conversions in a Single Calculation 172

■ Redefining the Kilogram 174

5.3 Mass Percent Composition 175

■ Iodized Salt 177

5.4 Using Mass Percent Composition to

Determine Empirical Formula 178

■ Fertilizer & Mass Percents 180

5.5 Using Empirical Formula and Molar Mass to Determine

Molecular Formula 181

viii

6 MOLECULAR SHAPE 192

6.1 Drawing Simple Lewis Structures 193

• Lewis Structures of Simple Molecules 193

• Lewis Structures of Molecules with a Central Atom 195 • Lewis Structures of Simple

Polyatomic Ions 195

6.2 Lewis Structures Continued 198

• Lewis Structures with Less Obvious Skeletal Structures 198 • Lewis Structures with Multiple Bonds 199 • Exceptions to the Octet Rule 200

■ Bleaching, Disinfecting, and Decontamination 200

6.3 Resonance Structures 201 6.4 Molecular Shape 203

■ Flavor, Molecular Shape, and Line Structures 204

• Bond Angles 208

■ Molecular Shapes Resulting from Expanded Octets 209

6.5 Electronegativity and Polarity 211

• Electronegativity 211 • Bond Polarity 213

• Intermolecular Forces in Review 224

CHANGES 234

7.1 General Properties of the Condensed Phases 235

7.2 Types of Solids 236

• Ionic Solids 236 • Molecular Solids 236

• Atomic Solids 238 • Network Solids 239

■ A Network Solid as Hard as Diamond 240

7.3 Physical Properties of Solids 243

• Vapor Pressure 243 • Melting Point 244

ix

■ Surface Tension and the Shape of Water Drops 248

• Vapor Pressure 249 • Boiling Point 250

■ High Altitude and High-Pressure Cooking 252

7.5 Energy and Physical Changes 253

• Temperature Changes 253 • Solid-Liquid Phase Changes: Melting and Freezing 255 • Liquid-Gas Phase Changes: Vaporization and Condensation 256 • Solid-Gas Phase Changes: Sublimation 257

8.3 The Gas Equations 277

• The Ideal Gas Equation 277

■ Pressure Exerted by a Column of Fluid 281

• The Combined Gas Equation 281

• The Molar Mass Gas Equation 282

8.4 The Gas Laws 285

• Boyle’s Law: The Pressure-Volume Relationship 285

• Charles’s Law: The Temperature-Volume Relationship 287

■ Automobile Air Bags and Charles’s Law 290

• Avogadro’s Law: The Moles-Volume Relationship 290

8.5 Gas Mixtures 292

• Dalton’s Law of Partial Pressures 292 • Mole Fractions 294

x

9 PHYSICAL PROPERTIES OF SOLUTIONS 308

9.1 General Properties of Solutions 309

■ Honey – A Supersaturated Solution 310

■ Instant Hot Packs 311

9.2 Aqueous Solubility 311 9.3 Solution Concentration 312

• Metals 349 • Nonmetals 349

• Noble Gases 349 • Metalloids 349

10.3 Balancing Chemical Equations 350

■ The Stoichiometry of Metabolism 354

10.4 Types of Chemical Reactions 355

■ Dental Pain and Redox 370

10.5 Chemical Reactions and Energy 372 10.6 Chemical Reactions in Review 372

xi

11 USING BALANCED CHEMICAL EQUATIONS 382

11.1 Mole to Mole Conversions 383

11.2 Mass to Mass Conversions 386

11.3 Limitations on Reaction Yield 387

• Limiting Reactant 388 • Percent Yield 391

■ Combustion Analysis 392

■ Alka-Seltzer 393

11.4 Aqueous Reactions 395

11.5 Gases in Chemical Reactions 400

• Predicting the Volume of a Gaseous Product 400 • Calculating the Required Volume of a Gaseous Reactant 401

■ Joseph Louis Gay-Lussac 403

11.6 Chemical Reactions and Heat 404

12.1 Properties of Acids and Bases 417

■ James Lind 418

12.2 Definitions of Acids and Bases 419

• Arrhenius Acids and Bases 419

• Brønsted Acids and Bases 419

• Conjugate Acid-Base Pairs 420

12.3 Water as an Acid; Water as a Base 422

12.4 Strong Acids and Bases 424

12.5 pH and pOH Scales 427

■ Antacids and the pH Balance in Your Stomach 434

13 EQUILIBRIUM 454

13.1 Reaction Rates 455

Collision Theory 458

13.2 Chemical Equilibrium 460

Reverse Processes Are Ongoing in a System

at Equilibrium? 462

13.3 Equilibrium Constants 462

■ Sweet Tea 463

• Calculating Equilibrium Constants 463

• Magnitude of the Equilibrium Constant 466

13.4 Factors that Affect Equilibrium 467

■ Hemoglobin Production at High Altitude 467

• Addition or Removal of a Substance 468

• Changes in Volume 470 • Changes in Temperature 471

■ Partially Hydrogenated Vegetable Oils 487

■ Representing Organic Molecules with Bond-Line Structures 489

14.4 Functional Groups 490 14.5 Alcohols and Ethers 491 14.6 Aldehydes and Ketones 493 14.7 Carboxylic Acids and Esters 495 14.8 Amines and Amides 496

14.9 Polymers 498

xiii

15 BIOCHEMISTRY 506

15.1 Biologically Important Molecules 507

• Glycerol 507 • Fatty Acids 507 • Amino Acids 507 • Sugars 508 • Phosphates 509

16.3 Dating Using Radioactive Decay 528

16.4 Medical Applications of Radioactivity 530

■ How Nuclear Chemistry Is Used to Treat Cancer 531

16.5 Nuclear Fission and Nuclear Fusion 531

Nuclear Fission and Fusion 532

xiv

17 ELECTROCHEMISTRY 538

17.1 Balancing Oxidation-Reduction Reactions Using the Half-Reaction Method 539 17.2 Batteries 543

Construction of a Galvanic Cell 544

• Dry Cells and Alkaline Batteries 547

• Lead Storage Batteries 548

• Lithium-Ion Batteries 549 • Fuel Cells 549

17.3 Corrosion 550 17.4 Electrolysis 552

• Electrolysis of Molten Sodium Chloride 552

From its very origin, Introductory Chemistry: An Atoms First Approach by Julia Burdge

and Michelle Driessen has been developed and written using an atoms first approach

specific to introductory chemistry It is not just a pared down version of a general istry text, but carefully crafted with the introductory-chemistry student in mind

chem-The ordering of topics facilitates the conceptual development of chemistry for the novice, rather than the historical development that has been used traditionally Its language and style are student friendly and conversational; and the importance and wonder of chemistry in everyday life are emphasized at every opportunity Continuing in the Burdge tradition, this text employs an outstanding art program, a consistent problem-solving approach, interesting applications woven throughout the chapters, and a wide range of end-of-chapter problems

Features

∙ Logical atoms first approach, building first an understanding of atomic structure,

followed by a logical progression of atomic properties, periodic trends, and how pounds arise as a consequence of atomic properties Following that, physical and chemical properties of compounds and chemical reactions are covered—built upon a solid foundation of how all such properties and processes are the consequence of the nature and behavior of atoms

com-∙ Engaging real-life examples and applications Each chapter contains relevant,

inter-esting stories in Familiar Chemistry segments that illustrate the importance of try to other fields of study, and how the current material applies to everyday life Many chapters also contain brief historical profiles of some important people in chemistry and other fields of scientific endeavor

chemis-Preface

278 CHAPTER 8 Gases

SAMPLE PROBLEM 8.2

Calculate the volume of a mole of ideal gas at room temperature (25°C) and 1.00 atm.

Strategy Convert the temperature in °C to temperature in kelvins, and use the ideal gas equation to solve for the unknown volume.

Setup The data given are n = 1.00 mol, T = 298 K, and P = 1.00 atm Because the pressure is expressed in atmospheres, we

use R = 0.0821 L · atm/K · mol to solve for volume in liters.

Practice Problem A TTEMPT What is the volume of 5.12 moles of an ideal gas at 32°C and 1.00 atm?

Practice Problem B UILD At what temperature (in °C) would 1 mole of ideal gas occupy 50.0 L (P = 1.00 atm)?

Practice Problem C ONCEPTUALIZE The diagram on the left represents a sample of gas in a container with a movable

piston Which of the other diagrams [(i)–(iv)] best represents the sample (a) after the absolute temperature has been doubled;

(b) after the volume has been decreased by half; and (c) after the external pressure has been doubled? (In each case, assume

that the only variable that has changed is the one specified.)

THINK ABOUT IT

With the pressure held constant, we should expect the volume to increase with increased temperature Room temperature

is higher than the standard temperature for gases (0 °C), so the molar volume at room temperature (25°C) should be higher

than the molar volume at 0 °C—and it is.

Using the Ideal Gas Equation to Calculate Volume

Student Note: It is a very common mistake to fail to convert to

absolute temperature when solving a gas problem Most often, temperatures are given in degrees Celsius The ideal gas equation only works when the temperature used is in kelvins

Remember: K = °C + 273.

(i) (ii) (iii) (iv)

SAMPLE PROBLEM 8.3

Calculate the pressure of 1.44 moles of an ideal gas in a 5.00-L container at 36°C.

Strategy Rearrange the ideal gas law (Equation 8.1) to isolate pressure, P Convert the temperature into kelvins, 36 + 273 = 309 K.

Using the Ideal Gas Equation to Calculate Pressure

∙ Consistent problem-solving skill

develop-ment Fostering a consistent approach to

problem solving helps students learn how

to approach, analyze, and solve problems

Each worked example (Sample Problem) is divided into logical steps: Strategy, Setup, Solution, and Think About It; and each is followed by three practice problems Prac-tice Problem A allows the student to solve

a problem similar to the Sample Problem, using the same strategy and steps Wherever possible, Practice Problem B probes under-standing of the same concept(s) as the Sam-ple Problem and Practice Problem A, but is sufficiently different that it requires a slightly different approach Practice Problem C often uses concept art or molecular models, and probes comprehension of underlying con-cepts The consistent use of this approach gives students the best chance for developing

a robust set of problem-solving skills

xvi

Molecular polarity is tremendously important in determining the physical and chemical properties of a substance Indeed,

we use a stepwise procedure:

1 Draw a correct Lewis structure [ Sections 6.1 and 6.2]

2 Count electron groups on the central atom Remember that an electron group can be a lone pair or a bond, and that

a bond may be a single bond, a double bond, or a triple bond.

3 Apply the VSEPR model [ Section 6.4] to determine electron-group geometry.

4 Consider the positions of the atoms to determine the molecular shape, which may or may not be the same as the electron-group geometry.

Consider the examples of SO 2 , C 2 H 2 , and CH 2 Cl 2 We determine the molecular shape of each as follows:

2 electron groups on each central atom:

Draw the Lewis structure

4 electron groups:

• 4 single bonds

With no lone pairs

on the central atom, the molecular shape is linear.

With 1 lone pair on the central atom, the molecular Consider positions

of atoms to determine molecular shape.

With no lone pairs

on the central atom, the molecular shape is tetrahedral.

2 electron groups arrange themselves linearly.

3 electron groups arrange themselves

in a trigonal plane.

Apply VSEPR to determine electron- group geometry

4 electron groups arrange themselves

in a tetrahedron.

Cl H C

Molecular Shape and Polarity KEY SKILLS

S and O have electronegativity values of 2.5 and Therefore, the bonds are polar.

Determine whether

or not the individual bonds are polar.

The C H bonds are nonpolar C and

Cl have electronegativity values of 2.5 and Therefore, the C Cl bonds are polar.

C

Cl Cl H S

Only in C 2 H 2 do the dipole-moment vectors cancel each other C 2 H 2 is nonpolar, SO 2 and CH 2 Cl 2 are polar.

Even with polar bonds, a molecule may be nonpolar if it consists of equivalent bonds that are distributed symmetrically

Molecules with equivalent bonds that are not distributed symmetrically—or with bonds that are not equivalent, even if they

are distributed symmetrically—are generally polar.

6.1 Determine the molecular shape of selenium dibromide.

a) linear b) bent c) trigonal planar d) trigonal pyramidal e) tetrahedral 6.2 Determine the molecular shape of phosphorus triiodide.

a) linear b) bent c) trigonal planar d) trigonal pyramidal e) tetrahedral

6.3 Which of the following species is polar?

∙ Outstanding pedagogy for student learning The Checkpoints and Student Notes

throughout each chapter are designed to foster frequent self-assessment and to provide timely information regarding common pitfalls, reminders of important information, and alternative approaches Rewind and Fast Forward Buttons help to illustrate and reinforce connections between material in different chapters, and enable students to find pertinent review material easily, when necessary

∙ Key Skills pages are reviews of specific skills that the authors know will be important

to students’ understanding of later chapters These go beyond simple reviews and ally preview the importance of the skills in later chapters They are additional opportu-nities for self-assessment and are meant to be revisited when the specific skills are required later in the book

actu-∙ Author-created online homework All of the online homework problems were

devel-oped entirely by co-author Michelle Driessen to ensure seamless integration with the book’s content

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Cindy Jolly Harwood, Purdue UniversityLindsay M Hinkle, Harvard University David G Jones, Vistamar SchoolBarbara S Pappas, The Ohio State UniversityAdditionally, we wish to thank our incredible team: Managing Director Thomas Timp, Director of Chemistry David Spurgeon, Director of Marketing Tami Hodge, Product Developer Robin Reed, Program Manager Lora Neyens, Content Project Manager Sherry Kane, Senior Designer David Hash, and Accuracy Checker John Murdzek

Julia Burdge and Michelle Driessen

Credits

Page iv (Julia Burdge): © David Spurgeon; p iv (Michelle Driessen): Courtesy of Michelle Driessen; Chapter 1:

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Model of the Atom

1.6 Isotopes

Atoms and Elements

The brilliant colors of a fireworks display result from the properties of the atoms

they contain These atoms give off specific colors when they are burned

Credit: © Jung-Pang Wu/Getty Images

3

In This Chapter, You Will Learn

Some of what chemistry is and how it is studied using the scientific method You will learn about atomic structure and you will become acquainted with the periodic table, how it is organized, and some

of the information it embodies

Things To Review Before You Begin

• Basic algebra

Have you ever wondered how an automobile airbag works? Or why iron rusts when exposed to water and air, but gold does not? Or why cookies “rise” as they bake? Or what causes the brilliant colors of fireworks displays? These phenomena, and countless

others, can be explained by an understanding of the fundamental principles of chemistry

Whether or not we realize it, chemistry is important in every aspect of our lives In the course of this book, you will come to understand the chemical principles responsible for many familiar observations and experiences

turn, is anything that has mass and occupies space Mass is one of the ways that

sci-entists measure the amount of matter.

You may already be familiar with some of the terms used in chemistry—even if

you have never taken a chemistry class You have probably heard of molecules; and even if you don’t know exactly what a chemical formula is, you undoubtedly know that

“H2O” is water You may have used or at least heard the term chemical reaction; and you are certainly familiar with many processes that are chemical reactions.

Why Learn Chemistry?

Chances are good that you are using this book for a chemistry class you are required

to take—even though you may not be a chemistry major Chemistry is a required part

of many degree programs because of its importance in a wide variety of scientific disciplines It sometimes is called the “central science” because knowledge of chemis-try supports the understanding of other scientific fields—including physics, biology, geology, ecology, oceanography, climatology, and medicine Whether this is the first in

a series of chemistry classes you will take or the only chemistry class you will ever take, we hope that it will help you to appreciate the beauty of chemistry—and to understand its importance in our daily lives

The Scientific Method

Scientific experiments are the key to advancing our understanding of chemistry or any science Although different scientists may take different approaches to experimentation,

we all follow a set of guidelines known as the scientific method This helps ensure the

quality and integrity of new findings that are added to the body of knowledge within

a given field

The scientific method starts with the collection of data from careful observations and/or experiments Scientists study the data and try to identify patterns When a pat-

tern is found, an attempt is made to describe it with a scientific law In this context,

a law is simply a concise statement of the observed pattern Scientists may then

for-mulate a hypothesis, an attempt to explain their observations Experiments are then

designed to test the hypothesis If the experiments reveal that the hypothesis is

incor-rect, the scientists must go back to the drawing board and come up with a different

interpretation of their data, and formulate a new hypothesis The new hypothesis will

then be tested by experiment When a hypothesis stands the test of extensive

experi-mentation, it may evolve into a scientific theory or model A theory or model is a

unifying principle that explains a body of experimental observations and the law or laws that are based on them Theories are used both to explain past observations and

to predict future observations When a theory fails to predict correctly, it must be

discarded or modified to become consistent with experimental observations Thus, by their very nature, scientific theories must be subject to change in the face of new data that do not support them

One of the most compelling examples of the scientific method is the development

of the vaccine for smallpox, a viral disease responsible for an estimated half a billion

deaths during the twentieth century alone Late in the eighteenth century, English cian Edward Jenner observed that even during smallpox outbreaks in Europe, a particu-

physi-lar group of people, milkmaids, seemed not to contract it.

Law: Milkmaids are not vulnerable to the virus that causes smallpox.

Based on his observations, Jenner proposed that perhaps milkmaids, who often

contracted cowpox, a similar but far less deadly virus from the cows they worked with,

had developed a natural immunity to smallpox

Hypothesis: Exposure to the cowpox virus causes the development of immunity

to the smallpox virus.

Jenner tested his hypothesis by injecting a healthy child with the cowpox virus—

and later with the smallpox virus If his hypothesis were correct, the child would not

contract smallpox—and in fact the child did not contract smallpox.

Theory: Because the child did not develop smallpox, immunity seemed to have resulted from exposure to cowpox.

Further experiments on many more people (mostly children and prisoners) firmed that exposure to the cowpox virus imparted immunity to the smallpox virus

con-The flowchart in Figure 1.1 illustrates the scientific method and how it guided the development of the smallpox vaccine

Observation:

Milkmaids don't

contract smallpox.

Further Experiment:

Many more humans inoculated with cowpox virus, confirming the model.

Hypothesis:

Having contracted

cowpox, milkmaids have a natural immunity

Experiment

Procedure to test hypothesis; measures one variable at a time

Model (Theory)

Set of conceptual assumptions that explains data from accumulated experiments;

predicts related phenomena

Further Experiment

Tests predictions based on model

Hypothesis revised if experimental results

do not support it

Model altered if experimental results

do not support it

Figure 1.1 Flowchart of the scientific method and its importance to Edward Jenner’s development of the smallpox vaccine.

atom is the smallest quantity of matter that still retains the properties of matter Further,

an element is a substance that cannot be broken down into simpler substances by any

means Common examples of elements include aluminum, which we all have in our kitchens in the form of foil; carbon, which exists in several different familiar forms—

including diamond and graphite (pencil “lead”); and helium, which can be used to fill

balloons The element aluminum consists entirely of aluminum atoms; the element carbon consists entirely of carbon atoms; and the element helium consists entirely of

helium atoms Although we can separate a sample of any element into smaller samples

of that element, we cannot separate it into other substances

Let’s consider the example of helium If we were to divide the helium in a balloon

in half, and then divide one of the halves in half, and so on, we would eventually (after a very large number of these hypothetical divisions) be left with a sample of helium consist-ing of just one helium atom This atom could not be further divided to give

two smaller samples of helium If this is difficult to imagine, think of a lection of eight identical iPods We could divide the collection in half three

col-times before we were left with a single iPod Although we could divide the

last iPod in half, neither of the resulting pieces would be an iPod! (Figure 1.2)The notion that matter consists of tiny, indivisible pieces has been around for a very long time, first having been proposed by the philosopher Democritus in the fifth century b.c But it was first formalized early in the nineteenth century by John Dalton (Figure 1.3) Dalton devised a theory to explain some of the most important observations made by scientists in the eighteenth century His theory included three statements, the first of which is:

∙ Matter is composed of tiny, indivisible particles called atoms; all atoms of a given element are identical; and atoms of one element are different from atoms of any other element

We will revisit this statement later in this chapter and introduce the second and third statements to complete our understanding of Dalton’s theory in Chapters 3 and 10

We know now that atoms, although very small, are not indivisible Rather,

they are made up of still smaller subatomic particles The type, number, and

arrangement of subatomic particles determine the properties of atoms, which in turn determine the properties of everything we see, touch, smell, and taste

Our goal in this book will be to understand how the nature of atoms gives rise to the properties of everything material To accomplish this, we will take a somewhat unconventional approach Rather than beginning with observations on the macroscopic scale and working our way backward to the atomic level of matter to explain these observations, we start by examin-ing the structure of atoms, and the nature and arrangement of the tiny subatomic particles that atoms contain

Student Note: By contrast,

consider a sample of salt water

We could divide it into smaller samples of salt water; but given the necessary equipment, we could also separate it into two different substances: water and salt An element is different in that it is not made up of other substances Elements are the simplest substances.

Figure 1.2 Repeatedly dividing this collection of iPods into smaller and smaller collections eventually leaves us with a single iPod, which we cannot divide further without destroying it.

for-a pfor-articulfor-ar type of color blindness, from which

he suffered This form of color blindness, where red and green cannot be distinguished, is known as Daltonism.

Credit: © Sheila Terry/Science Source

Before we begin our study of atoms, it is important for you to understand a bit about the behavior of electrically charged objects We are all at least casually familiar with the concept of electric charge You may have brushed your hair in very low humidity and had it stand on end; and you have certainly experienced static shocks and seen lightning All of these phenomena result from the interactions of electric charges The following list illustrates some of the important aspects of elec-tric charge:

∙ An object that is electrically charged may have a positive (+) charge or a negative (−) charge

positive

1

negative

2

∙ Objects with opposite charges (one negative and one positive) are attracted

to each other (You’ve heard the adage “opposites attract.”)

Nuclear Model of the Atom

Experiments conducted late in the nineteenth century indicated that atoms, which had

been considered the smallest possible pieces of matter, contained even smaller particles

The first of these experiments were done by J J Thomson, an English physicist The experiments revealed that a wide variety of different materials could all be made to

emit a stream of tiny, negatively charged particles—that we now know as electrons

Thomson reasoned that because all atoms appeared to contain these negative particles

but were themselves electrically neutral, they must also contain something positively

Credit: © Lourens Smak/Alamy

Credit: © Michael ONeill/WeatherVideoHD.TV

charged This gave rise to a model of the atom as a sphere of positive charge, throughout which negatively charged electrons were uniformly dis-tributed (Figure 1.4) This model was known as the “plum-pudding”

model—named after a then-popular English dessert Thomson’s pudding model, which was generally accepted for a number of years, was

plum-an early attempt to describe the internal structure of atoms Although it was generally accepted for a number of years, this model ultimately was proven wrong by subsequent experiments

Working with Thomson, New Zealand physicist Ernest Rutherford (one of Thomson’s own students) devised an experiment to test the plum-pudding model of atomic structure By that time, Rutherford had already

established the existence of another subatomic particle known as an alpha particle,

which is emitted by some radioactive substances Alpha particles are positively charged,

and are thousands of times more massive than electrons In his most famous experiment, Rutherford directed a stream of alpha particles at a thin gold foil A schematic of the experimental setup is shown in Figure 1.5 If Thomson’s model of the atom were cor-rect, nearly all of the alpha particles would pass directly through the foil—although a small number would be deflected slightly by virtue of passing very close to electrons

Rutherford surrounded the gold foil target with a detector that produced a tiny flash of light each time an alpha particle collided with it This allowed Rutherford to determine the paths taken by alpha particles Figure 1.6 illustrates the expected experimental result

The actual experimental result was very different from what had been expected

Although most of the alpha particles did pass directly through the gold foil, some were deflected at much larger angles than had been anticipated Some even bounced off the foil back toward the source—a result that Rutherford found absolutely shocking He knew that alpha particles could only be deflected at such large angles, and occasionally bounce back in the direction of their source, if they encountered something within the gold atoms that was (1) positively charged, and (2) much larger than themselves

Figure 1.7 illustrates the actual result of Rutherford’s experiment

This experimental result gave rise to a new model of the internal structure of atoms Rutherford proposed that atoms are mostly empty space, but that each has a

tiny, dense core that contains all of its positive charge and nearly all of its mass This

core is called the atomic nucleus.

Positively charged sphere

1 2

2 2

2 2 2 2 2 2

1 1 1 1

1 1 1 1

Figure 1.4 Thomson’s experiments indicated that atoms contained negatively charged particles, which he envisioned as uniformly distributed in a sphere of positive charge.

α particle emitter

Gold foil

Zinc-sulfide screen

Light flashes produced

by α particles hitting screen

Figure 1.5 Rutherford’s experiment directed a stream of positively charged alpha particles at a gold foil The nearly circular detector emitted a flash of light when struck by an alpha particle.

Electrons: tiny, negatively charged particles, uniformly distributed throughout the sphere

Path followed by alpha particles, directed at the gold foil

Gold atom: sphere of uniform positive charge

Figure 1.6 Rutherford’s gold foil experiment was designed to test Thomson’s plum-pudding model of the atom, which depicted the atom as negatively charged electrons uniformly distributed

in a sphere of positive charge If the model had been correct, the alpha particles would have passed directly through the foil, with a few being deflected slightly by interaction with electrons

(Remember that a positively charged object and a negatively charged object are attracted to each other A positively charged alpha particle could be pulled slightly off course if it passed very close to one of the negatively charged electrons.)

Subsequent experiments supported Rutherford’s nuclear model of the atom; and

we now know that all atomic nuclei (the plural of nucleus) contain positively charged

particles called protons And with the exception of hydrogen, the lightest element, atomic nuclei also contain electrically neutral particles called neutrons Together, the

protons and neutrons in an atom account for nearly all of its mass, but only a tiny tion of its volume The nucleus is surrounded by a “cloud” of electrons—and just as Rutherford proposed, atoms are mostly empty space Figure 1.8 illustrates the nuclear model of the atom

frac-Of the three subatomic particles in our model of the atom, the electron is the smallest and lightest Protons and neutrons have very similar masses, and each is nearly

Path followed by alpha particles directed at the gold foil

Gold nucleus: tiny, dense, positively charged center

Figure 1.7 The actual result of Rutherford’s gold foil experiment Positively charged alpha particles were directed at a gold foil Most passed through undeflected, but a few were deflected

at angles much greater than expected—some even bounced back toward the source This indicated that as they passed through the gold atoms, they encountered something positively charged and significantly more massive than themselves.

Nucleus containing protons (   ) and neutrons (   ) 

2

2 2

Figure 1.8 Nuclear model of the atom Protons (blue) and neutrons (red) are contained within the nucleus, a tiny space at the center of the atom The rest of the volume of the atom is nearly empty, but is occupied by the atom’s electrons This illustration exaggerates the size of the nucleus relative to the size of the atom If the picture were actually done to scale, and the nucleus were the size shown here (1 centimeter), the atom would be on the order of 100 meters across—about the length of a football field.

Student Note: An alpha particle

is the combination of two protons

and two neutrons.

2000 times as heavy as an electron Further, because protons are positively charged and electrons are negatively charged, combination of equal numbers of each results in com-plete cancellation of the charges The number of electrons is equal to the number of protons in a neutral atom Because neutrons are electrically neutral, they do not con-tribute to an atom’s overall charge

Sample Problem 1.1 lets you practice identifying which combinations of atomic particles constitute a neutral atom

sub-SAMPLE PROBLEM 1.1 Identifying a Neutral Atom Using Numbers of Subatomic Particles

The following table contains data sets that indicate numbers of subatomic particles Which of the sets of data represent neutral atoms? For those that do not represent neutral atoms, determine what the charge is—based on the numbers of subatomic particles.

neutrons protons electrons

(a) 5 10 5 (b) 11 12 12 (c) 8 9 9 (d) 20 21 20

Strategy You have learned that the charge on a proton is +1 and the charge on an electron is −1 Neutrons have no charge The overall charge is the sum of charges of the protons and electrons, and a neutral atom has no charge Therefore, a set of data in which the number of protons is equal to the number of electrons represents a neutral atom.

Setup Data sets (b) and (c) each contain equal numbers of protons and electrons Data sets (a) and (d) do not.

Solution The data in sets (b) and (c) represent neutral atoms Those in (a) and (d) represent charges species The charge on the species represented by data set (a) is +5: 10 protons (+1 each) and 5 electrons (−1 each) The charge on the species represented

by data set (d) is +1: 21 protons (+1 each) and 20 electrons (−1 each).

Practice Problem A TTEMPT Which of the following data sets represent neutral atoms? For those that do not represent neutral atoms, determine the charge.

neutrons protons electrons

(a) 31 31 30 (b) 24 22 24 (c) 12 11 11

Practice Problem B UILD Fill in the appropriate missing numbers in the following table:

overall charge protons electrons

1.4 Elements and the Periodic Table

The identity of an element is determined by the number of protons in its nucleus For example, an atom with two protons in its nucleus is helium; one with six protons is carbon;

and one with 79 protons is gold There are no helium atoms with any number other than two protons, no carbon atoms with any number other than six protons, and no gold atoms with any number other than 79 protons The number of protons in an atom’s nucleus is also

known as the atomic number, for which we use the symbol Z All of the known elements are arranged in order of increasing atomic number on the periodic table (Figure 1.9).

1.3.1 Which of the following can change without changing

the charge on an atom? (Select all that apply.) a) Number of protons c) Number of electrons b) Number of neutrons

1.3.2 Which of the following can change without changing the

elemental identity of an atom? (Select all that apply.) a) Number of protons c) Number of electrons b) Number of neutrons

1.3.3 Which of the following must be equal for the

combina-tion to constitute a neutral atom?

a) Number of protons and number of neutrons b) Number of protons and number of electrons c) Number of neutrons and number of electrons d) Number of protons, number of neutrons, and number

of electrons

1.3.4 Which of the following could represent a neutral atom?

(Select all that apply.)

Figure 1.9 The modern

periodic table The

elements are arranged

in order of increasing

atomic number, which

is shown above each

1A 1 2A 2

Sample Problem 1.2 lets you practice identifying an element, given its atomic number

The extraordinary abundance of oxygen results from our bodies containing so much water (89 percent of water’s mass is the oxygen it contains)

Depending on health and age, the water content of a human body can range from 50 percent in a dehydrated person to 75 percent in a healthy infant.

The second most abundant element in our bodies, carbon, actually has a fairly low natural abundance Although it makes up only about 0.1 percent

of Earth’s crust, carbon is present in nearly all living systems.

Elements in the Human Body

Although the human body contains trace amounts of a large variety of elements, nearly 99 percent

of our mass consists of just six of the 118 known elements:

Identify the element given the atomic number of 16.

Strategy You have learned that the atomic number of an element represents the number of protons the element contains The atomic number is found above the element’s symbol on the periodic table.

Setup The element contains 16 protons.

Solution The elements on the periodic table are arranged in order of increasing atomic number The symbol on the periodic table that has a 16 above its symbol is S This symbol represents the element sulfur.

Practice Problem A TTEMPT Identify the element with an atomic number of 35.

Practice Problem B UILD Determine the atomic number for iodine.

Practice Problem C ONCEPTUALIZE Identify the atomic number and identity of each atom shown (Protons are blue, neutrons are red, and electrons are green.)

The periodic table also identifies each element with a chemical symbol A

chemi-cal symbol consists of one capital letter, or a combination of two letters, one capital and one lowercase The chemical symbol for helium, for example, is He, and that for carbon is C Most chemical symbols, including He and C, are derived from the familiar English names of the elements

Others are derived from an element’s Greek or Latin name and may take some

practice for you to recognize Examples include Au (aurium) for gold, Sn (stannum) for tin, Na (natrium) for sodium, and K (kalium) for potassium Many of the most recent

additions to the periodic table (the highest atomic numbers) are named to honor the scientists involved in their discovery

Spend some time looking at the periodic table shown in Figure 1.9, or on the inside front cover of this book Note that each square on the table contains a chemical symbol and a number, along with the element’s name The number at the top of each

square is the atomic number, which is always a whole number (Remember that the atomic number, Z, is the number of protons.) Each element can be identified by its

atomic number, its name, or its chemical symbol—and we need only one of these pieces

of information to unambiguously specify the identity of an element The periodic table squares for helium, carbon, and gold are:

Gold

Au79

HeliumHe2

CarbonC6

Sample Problem 1.3 lets you practice using atomic number, name, and chemical

symbol to identify an element

Complete the following table:

Element Chemical Symbol Atomic Number

(a) calcium

Strategy You have learned that the atomic number of an element represents the number of protons the element contains and is

found above the element’s symbol on the periodic table.

Setup Using the one given piece of information in each part, the other two can be found If the chemical name is given, the

symbol should be determined and used to find the atomic number using the periodic table If the chemical symbol is given, it

should be used to determine the name of the element and the atomic number shown on the periodic table If the atomic number

is given, it should be found on the periodic table to determine the chemical symbol and element name.

Solution In part (a) the chemical symbol for calcium is Ca Using the periodic table, locate Ca and find that its atomic number

is 20 Part (b) gives the chemical symbol for copper The chemical symbol Cu can be located on the periodic table to determine

the atomic number is 29 Part (c) gives the atomic number, which can be located on the periodic table to find that the chemical

symbol for the element is Al This symbol represents the element aluminum.

THINK ABOUT IT

A strong grasp of the names and chemical symbols of the elements will allow you to use the periodic table to determine

many properties of any element in question and vice versa.

Relating an Element’s Identity to Its Chemical Symbol and Atomic Number

Practice Problem A TTEMPT Complete the following table:

Element Chemical Symbol Atomic Number

(a) rubidium

Practice Problem B UILD Identify the sets of data that are incorrect in the table:

Element Chemical Symbol Atomic Number

(b) strontium Sr 38

Practice Problem C ONCEPTUALIZE Complete the following table:

Element Chemical Symbol Atomic Number (Protons) Neutrons Electrons

1.4.2 Which pieces of information are sufficient for you to identify an element? (Select all that apply.)

a) Atomic number only b) Atomic number and element symbol c) Elemental symbol only

d) Element name only e) Element name and element symbol

Credit: © ericsphotography/iStock/Getty Images

Familiar Chemistry

Helium

We have all seen helium balloons used as decorations and gifts; and most of us have been entertained

by the silly-sounding high-pitched voice of a person who has breathed in the helium from a balloon

But as familiar as this may be, how much do you really know about helium? Where does it come from?

How abundant is it? Why does a balloon filled with helium float in the air? And what other uses do

we have for the element helium? Helium is actually the product of a radioactive decay process, and although you may not understand yet what that is, you are probably aware that uranium is

“radioactive.” As it turns out, part of what makes uranium radioactive is the process that produces helium On Earth, helium is found in and around natural gas deposits, and although it is relatively rare here on Earth, it is the second most abundant element in the universe The element helium was discovered late in the nineteenth century—and its value to society has been immense It

is indispensable as coolant for magnetic resonance imaging (MRI) machines; it is used in the manufacture of computer chips, in scuba diving gas mixtures, in arc welding operations, and in a host of military applications—including air-to-air missile guidance and surveillance operations Helium balloons float because helium is “lighter” than air (Technically, helium has a lower density [ Section 4.4]

than air.) It is precisely because helium rises that we are facing a shortage here on Earth Helium that

is released into the air will rise until it leaves the atmosphere and floats out into space Helium is considered a nonrenewable resource, prompting large-scale users of it (the military, the medical industry, scientific research facilities, and the silicon-chip industry) to develop methods for capturing and recycling the helium that they use.

1.5 Organization of the Periodic Table

The periodic table (Figure 1.9) consists of 118 elements, arranged in vertical columns

called groups and horizontal rows called periods You will notice that the groups are

headed by numerical designations The top designation, comprising a number and a letter,

is the most commonly used system in which the table is divided into main-group elements (designated A), and transition elements (designated B) The main-group elements include

Groups 1A and 2A on the left, and 3A through 8A on the right The transition elements

are those in the sunken, middle section of the table, with B group designations (The groups

are also numbered 1 through 18 from left to right However, throughout this book, we will use the A and B numbering system consistently to refer to groups of the periodic table.)

Although the periodic table is now arranged in order of increasing atomic number

(left to right, starting at the top), it was arranged originally in groups of elements with similar properties—even before the concept of atomic number was known Thus, the proper-ties of elements within a group tend to be similar Some of the groups have special names that refer to the shared properties of the elements they contain Group 1A, for example, is

called the alkali metals; Group 2A is called the alkaline earth metals; Group 6A is called the chalcogens; Group 7A is called the halogens; and Group 8A is called the noble gases.

In addition to groups (columns) and periods (rows), the periodic table is divided

into metals and nonmetals by the diagonal zigzag line on the right side of the table Most

elements are metals (left of the zigzag line) Nonmetals are to the right of the zigzag line

A handful of elements have properties that are intermediate between metal and nonmetal

and are referred to as metalloids These are found adjacent to the zigzag line By noting

an element’s position in the periodic table, you can determine whether it is a metal, a nonmetal, or a metalloid

Sample Problem 1.4 gives you some practice classifying elements by their positions

on the periodic table

Student Note: The properties

that distinguish metals and

nonmetals will be discussed in

Chapter 2 [ Section 2.6] ,

however, you are undoubtedly

familiar with the term metal and

have a good sense of what

constitutes metallic properties

Metals conduct electricity and

most are shiny solids.

Identify each of the following elements as a metal, nonmetal, or metalloid:

(a) N (b) Si (c) Ca (d) Cl (e) As

Strategy Find the given chemical symbol on the periodic table.

Setup The nonmetallic elements are found in the upper right corner of the periodic table, above the zigzag line The metallic

elements are found below the zigzag line Note that the metalloids include the highlighted symbols next to the zigzag line The

metalloids are neither metals nor nonmetals.

Solution Part (a) gives the symbol for nitrogen which is found in the nonmetal portion of the periodic table Part (b) describes

silicon, which is a metalloid found along the zigzag line Part (c) shows calcium, found in the metals area of the periodic table

Part (d) is chlorine, a nonmetal Part (e) describes arsenic, a metalloid.

THINK ABOUT IT

Most of the periodic table is composed of metals, with the elements in the upper right corner being nonmetals A tiny

number of elements are shaded along the zigzag line and are considered metalloids.

Identifying an Element as Metal, Nonmetal, or Metalloid by Its Position on the Periodic Table

Practice Problem A TTEMPT Identify each of the following elements as a metal, nonmetal, or metalloid:

(a) Se (b) Al (c) Na (d) Kr (e) Ge

Practice Problem B UILD Name an element that fits each of the following descriptions Note that there may be more than

one element that will work.

(a) A nonmetal found in group 14 (4A) (d) A nonmetal found in group 15 (5A).

(b) A metalloid found in group 13 (3A) (e) A metal found in group 14 (4A).

(c) A metal found in group 15 (5A).

Practice Problem C ONCEPTUALIZE Determine which categories each element (chemical symbol given) falls into

Rubidium is used as an example.

Symbol Main-Group Element Transition Element Metal Nonmetal Metalloid Alkali Metal Earth Metal Alkaline Halogen Noble Gas

B Zn K

Familiar Chemistry

Elements in Earth’s Crust

Earth’s crust extends from the surface to an average depth of about

40 km (25 mi) Of the 118 known elements, just eight elements make

up nearly 99 percent of our planet’s crust They are, in decreasing order of abundance, oxygen (O), silicon (Si), aluminum (Al), iron (Fe), calcium (Ca), sodium (Na), potassium (K), and magnesium (Mg)

Beneath the crust is the mantle, a hot, fluid mixture of iron, carbon (C), silicon, and sulfur (S); and a solid core believed to consist mostly

of iron.

Of the eight most abundant elements, oxygen and silicon alone constitute over 70 percent of the crust These two elements combine (along with small amounts of other elements) to form a huge variety

of silicate minerals, including the two most common minerals, feldspar and quartz The feldspar and quartz families of minerals include many familiar rocks and gemstones.

Core

MantleCrust

Oxygen47%

Aluminum 8.1%

Silicon28%

Credit: © John Cancalosi/Getty Images Credit: © Doug Sherman/Geofile Credit: © Harry Taylor/Getty Images

Feldspar minerals:

Quartz minerals:

Credit: © Doug Sherman/Geofile Credit: © Dr Parvinder Sethi Credit: © Ron Evans/Getty Images

1.5.1 Which of the following series of elemental symbols

lists a nonmetal, a metal, and a metalloid?

a) Ca, Cu, Si c) Br, Ba, Ge e) Ag, Cr, As b) K, Mg, B d) O, Na, S

1.5.2 Which of the following elements would you expect to

have properties most similar to those of chlorine (Cl)?

a) Cu c) Na e) S b) F d) Cr

1.5.3 The periodic table shown here has four groups highlighted Which of the highlighted groups contain(s) only one metal, which contain(s) only one nonmetal, and which contain(s) only one metalloid?

a) iv, iii, and i d) iii & iv, ii, and i b) i, iv, and iii e) i, ii, and iii & iv c) iv, ii, and iii

We have learned that an atom can be identified by the number of protons contained

in its nucleus—also known as its atomic number, Z But remember that with just one exception (hydrogen), the nuclei of atoms also contain neutrons—and most ele- ments consist of mixtures of atoms with different numbers of neutrons For example,

in a sample of pure chlorine, all of the atoms have 17 protons—but they do not all  have the same number of neutrons Roughly three-quarters of the chlorine atoms  will have 18 neutrons and one-quarter will have 20 neutrons An atom with

17 protons and 18 neutrons and an atom with 17 protons and 20 neutrons are both

chlorine atoms They are, however, different isotopes of chlorine Isotopes are

atoms of the same element, and therefore have the same number of protons, but have different numbers of neutrons

The mass number (A) is the total number of neutrons and protons in an atom’s nucleus (Protons and neutrons are known, collectively, as nucleons.) Returning to

the example of chlorine, the mass number of a chlorine atom with 18 neutrons is

35 (17 protons + 18 neutrons), and the mass number of a chlorine atom with 20 trons is 37 (17 protons + 20 neutrons) The way to denote the identity of an atom is

neu-Student Note: Note that this is

contrary to the part of Dalton’s

atomic theory that we have

encountered Atoms of a given

element are, in fact, not identical.

(number of protons)

There are three isotopes of hydrogen, called hydrogen, deuterium, and tritium

Hydro-gen has one proton and no neutrons in its nucleus, deuterium has one proton and one neutron, and tritium has one proton and two neutrons Thus, to represent the isotopes

of hydrogen, we write

1

1Hprotium

2

1Hdeuterium

3

1Htritium

Similarly, the two common isotopes of uranium (Z = 92), which have mass numbers

of 235 and 238, respectively, can be represented as follows:

235

92UThe first isotope, with 235 − 92 = 143 neutrons in its nucleus, is used in nuclear reac-tors and atomic bombs, whereas the second isotope, with 146 neutrons, lacks the prop-erties necessary for these applications With the exception of hydrogen, which has different names for each of its isotopes, the isotopes of other elements are identified by their mass numbers The two isotopes of uranium are called uranium-235 (pronounced

“uranium two thirty-five”) and uranium-238 (pronounced “uranium two thirty-eight”)

Because the atomic number subscript can be deduced from the elemental symbol, it may

be omitted from these representations without the loss of any information The symbols

3H and 235U are sufficient to specify the isotopes tritium and uranium-235, respectively

The chemical properties of an element are determined primarily by the number

of protons and electrons in its atoms, not by the number of neutrons Therefore, isotopes

of the same element typically exhibit very similar chemical properties

Sample Problem 1.5 shows how to calculate the number of protons, neutrons, and electrons using atomic numbers and mass numbers

a positively charged ion (with a certain charge-to-mass, e/m ratio)

The positive ions are accelerated between two plates of opposite charges The beam of ions then passes through a magnetic field, which separates the ions on the basis of the e/m ratio The smaller the e/m ratio, the more the magnetic field deflects the ion The magnitude

of deflection is used to determine the mass of the ion and therefore the mass of the parent atom.

Thinking Outside the Box

Filament Ion beam Magnet

Detecting screen

Sample gas

Accelerating plates Electron

beam

Diagram of one type of mass spectrometer.

Student Note: Because these

symbols designate isotopes by specifying numbers of nucleons, they are sometimes referred to

as nuclear symbols.

1.6.1 How many neutrons are there in an atom of 60 Ni?

a) 60 d) 32 b) 30 e) 29 c) 28

1.6.2 What is the mass number of an oxygen atom with nine

neutrons in its nucleus?

a) 8 d) 16 b) 9 e) 18 c) 17

Strategy Recall that the superscript denotes the mass number, A, and the subscript denotes the atomic number, Z In cases

where no subscript is shown, as in parts (b), (c), and (d), the atomic number can be deduced from the elemental symbol or name

For the purpose of determining the number of electrons, remember that atoms are neutral, so the number of electrons is equal to

the number of protons.

Setup Number of protons = Z, number of neutrons = A − Z, and number of electrons = number of protons Recall that the 14

in carbon-14 is the mass number.

Solution (a) The atomic number is 17, so there are 17 protons The mass number is 35, so the number of neutrons is 35 − 17 = 18

The number of electrons equals the number of protons, so there are 17 electrons.

(b) Because the element is again Cl (chlorine), the atomic number is again 17, so there are 17 protons The mass number is 37,

so the number of neutrons is 37 − 17 = 20 The number of electrons equals the number of protons, so there are 17 electrons, too.

(c) The atomic number of K (potassium) is 19, so there are 19 protons The mass number is 41, so there are 41 − 19 = 22

neutrons There are 19 electrons.

(d) Carbon-14 can also be represented as 14 C The atomic number of carbon is 6, so there are 6 protons and 6 electrons There

are 14 − 6 = 8 neutrons.

Determining Numbers of Subatomic Particles in a Given Atom

THINK ABOUT IT

Verify that the number of protons and the number of neutrons for each example sum to the mass number that is given In

part (a), for example, there are 17 protons and 18 neutrons, which sum to give a mass number of 35, the value given in the

problem In part (b), 17 protons + 20 neutrons = 37 In part (c), 19 protons + 22 neutrons = 41 In part (d), 6 protons +

8 neutrons = 14.

Practice Problem A TTEMPT How many protons, neutrons, and electrons are there in an atom of (a) 10

5 B, (b) 36 Ar, (c) 38 Sr, and (d) carbon-11?

Practice Problem B UILD Give the correct symbols to identify an atom that contains (a) 4 protons, 4 electrons, and

5 neutrons; (b) 23 protons, 23 electrons, and 28 neutrons; (c) 54 protons, 54 electrons, and 70 neutrons; and (d) 31 protons,

31 electrons, and 38 neutrons.

Practice Problem C ONCEPTUALIZE Fill in the missing information for neutral atoms:

Isotope Symbol Element Name Number (A)Mass Neutrons (n°) Protons (p + ) Electrons (e − )

nitrogen 14

As we have seen, there are two different isotopes of chlorine, 35Cl and 37Cl However,

if you examine the periodic table on the inside cover of the book, you will see a ber under the element’s symbol and name that is neither 35 nor 37 The number under

num-the symbol and name of chlorine is 35.45 That number, 35.45, is num-the atomic mass (M)

of chlorine In order to understand the concept of atomic mass, we need to become

familiar with the units with which atomic masses are expressed—namely the atomic

atom On this scale, the mass of a 35Cl atom is 34.968852721 amu (Note that the mass

of an atom is not exactly equal to its mass number.) The mass of a 37Cl atom is 36.96590262 amu

The number on the periodic table, 35.45, is the average atomic mass of

chlo-rine It is much closer to 35 than it is to 37, though, because it is not a simple average, but rather a weighted average Because 35Cl is far more abundant in naturally occur-ring chlorine than 37Cl, the average atomic mass of chlorine is closer to the mass of

Familiar Chemistry

Iron-Fortified Cereal

Iron deficiency is the single most common nutritional deficiency in the world An estimated 25 percent of the world’s population does not consume enough iron to maintain good health Iron is necessary for production of hemoglobin, the component in red blood cells responsible for the transport of oxygen, and insufficient iron causes anemia People with anemia can suffer from a variety of symptoms including fatigue, weakness, pale color, poor appetite, headache, and light-headedness Although pharmacy shelves display a variety of over-the-counter iron supplements, one of the most popular ways to increase dietary intake of iron is by eating iron-fortified cereal Such cereals are common and include many familiar brands Have you ever thought about how the cereals become “fortified”? It may surprise you to learn that most cereals are fortified with iron simply by the addition of iron metal! Finely divided bits of iron are simply added to the grain and other ingredients during processing The iron metal in fortified cereals is fairly simple to separate and observe—and this process is a popular chemistry demonstration If you blend the cereal with water and apply a strong magnet, you can actually separate the iron filings.

Credit: (both) © McGraw-Hill Education/David Moyer, photographer

(A number of YouTube videos illustrate this process quite nicely Search: iron fortified cereal)