Chapter 4 Structure of the Atom

CHAPTER

4

The Structure of the Atom

Resource Manager
Section

86A

Objectives

Activities/Features

Section 4.1

1. Compare and contrast the atomic

Discovery Lab: Observing Electrical

Early Theories of Matter
1 session
1/2 block

models of Democritus and Dalton.
2. Define an atom.

Charge, p. 87

Section 4.2

3. Distinguish between the subatomic

Problem-Solving Lab: Interpreting STM

Subatomic Particles and the Nuclear Atom
1 session
1/2 block

particles in terms of relative charge and mass.
4. Describe the structure of the nuclear atom, including the locations of the subatomic particles.

Images, p. 96

Section 4.3

5. Explain the role of atomic number in

MiniLab: Modeling Isotopes, p. 102

How Atoms Differ
2 sessions
1 block

determining the identity of an atom.
6. Define an isotope and explain why atomic masses are not whole numbers.
7. Calculate the number of electrons, protons, and neutrons in an atom given its mass number and atomic number.

Section 4.4

8. Explain the relationship between

Careers Using Chemistry: Radiation

Unstable Nuclei and
Radioactive Decay
2 sessions
1 block

unstable nuclei and radioactive decay.
9. Characterize alpha, beta, and gamma radiation in terms of mass and charge.

ChemLab: Very Small Particles, pp. 108–109
Chemistry and Society: Nanotechnology,

History Connection, p. 90

Protection Technician, p. 106

p. 110

CHAPTER 4 RESOURCE MANAGER

National Science
Content Standards

State/Local
Standards

Reproducible Masters

Transparencies

UCP.1, UCP.2; A.1; B.1,
B.2, B.4; G.2, G.3

Study Guide for Content Mastery,
p. 19 L2

Section Focus
Transparency 13 L1 ELL

UCP.1, UCP.2; A.2; B.1,
B.4; G.1, G.2, G.3

Study Guide for Content Mastery,
p. 20 L2
Laboratory Manual, pp. 25–28

Section Focus
Transparency 14 L1 ELL
Teaching
Transparencies 11, 12 L2 ELL

L2

UCP.1, UCP.2; A.1; B.1,
B.2; G.2

Study Guide for Content Mastery, pp. 21–23 L2
ChemLab and MiniLab
Worksheets, p. 13 L2
Challenge Problems, p. 4 L3

Section Focus
Transparency 15 L1 ELL
Teaching
Transparency 13 L2 ELL
Math Skills
Transparency 4 L2 ELL

UCP.1, UCP.2; A.1; B.1,
B.2, B.4, B.6; E.1, E.2;
F.5, F.6; G.3

Study Guide for Content Mastery,
p. 24 L2
ChemLab and MiniLab
Worksheets, pp. 14–16 L2
Laboratory Manual, pp. 29–32

Section Focus
Transparency 16 L1 ELL
Teaching
Transparency 14 L2 ELL

L2

Key to National Science Content Standards: UCP ϭ Unifying Concepts and
Processes, A ϭ Science as Inquiry, B ϭ Physical Science, C ϭ Life Science,
D ϭ Earth and Space Sciences, E ϭ Science and Technology,
F ϭ Science in Personal and Social Perspectives, G ϭ History and Nature of Science
Refer to pages 4T–5T of the Teacher Guide for an explanation of the
National Science Content Standards correlations.

86B

CHAPTER

4

The Structure of the Atom

Resource Manager
Materials List
ChemLab (pages 108–109) vanilla extract (6 mL), 9-inch latex balloon (2), dropper

Discovery Lab (page 87) plastic comb, hole punch, paper, ruler, clear plastic tape (four 10-cm long pieces)

MiniLab (page 102) balance, bag of pre- and post-1982 pennies

Demonstration (pages 92–93)
Crooke’s tube, high-voltage power supply, bar magnet

Preparation of Solutions
For a review of solution preparation, see page 46T of the Teacher Guide.
There are no solutions to be prepared for the activities in this chapter.

86C

Assessment Resources

Additional Resources

Chapter Assessment, pp. 19–24
MindJogger Videoquizzes DVD/VHS
Alternate Assessment in the Science Classroom
ExamView® Pro CD-ROM
Solutions Manual, Chapter 4
Supplemental Problems, Chapter 4
Performance Assessment in the Science Classroom
Chemistry Interactive CD-ROM, Chapter 4 quiz

Spanish Resources ELL
Guided Reading Audio Program, Chapter 4 ELL
Cooperative Learning in the Science Classroom
Lab and Safety Skills in the Science Classroom
Lesson Plans
Block Scheduling Lesson Plans
Chemistrymc.com

CHAPTER 4 RESOURCE MANAGER

Glencoe Technology
The following multimedia for this chapter are available from Glencoe.
VIDEOTAPE/DVD

CD-ROM

MindJogger Videoquizzes,
Chapter 4

ExamView® Pro
Testmaker

CD-ROM
Chemistry: Matter and Change
History of the Atomic Theory, Video
Discovery of the Electron, Experiment
Thomson’s Experiment, Demonstration
Rutherford’s Gold, Animation
Evidence for Alpha Particles,
Demonstration

CD-ROM

Interactive Chalkboard
CD-ROM

TeacherWorks™

Multiple Learning Styles
Look for the following icons for strategies that emphasize different learning modalities.
Kinesthetic
Intrapersonal
Differentiated Instruction, pp. 95, 98, 106;
Enrichment, p. 88; Extension, p. 89
Reteach, p. 97
Linguistic
Visual-Spatial
Chemistry Journal, pp. 88, 100; Reteach, p. 91;
Differentiated Instruction, p. 89; Chemistry
Portfolio, p. 101; Enrichment, p. 93; Differentiated
Journal, p. 94; Extension, p. 97; Enrichment, p. 103
Instruction, p. 94
Interpersonal
Logical-Mathematical
Check for Understanding, p. 107
Check for Understanding, p. 104

Key to Teaching Strategies
L1 Level 1 activities should be appropriate for students with learning difficulties.
L2 Level 2 activities should be within the ability range of all students.
L3 Level 3 activities are designed for above-average students.
ELL ELL activities should be within the ability range of
English Language Learners.
COOP LEARN Cooperative Learning activities are designed for small group work.
P These strategies represent student products that can be placed into a best-work portfolio.

Assessment Planner
Portfolio Assessment
Portfolio, TWE, pp. 101, 111
Assessment, p. 95
Performance Assessment
Assessment, TWE, pp. 91, 101
ChemLab, TWE, p. 109
Problem-Solving Lab, TWE,
p. 96
MiniLab, TWE, p. 102
ChemLab, SE, pp. 108–109
Discovery Lab, SE, p. 87
MiniLab, SE, p. 102

Knowledge Assessment
Assessment, TWE, p. 107
Section Assessment, SE, pp. 91, 97, 104, 107
Chapter Assessment, SE, pp. 112–115
Skill Assessment
Assessment, TWE, pp. 89, 97,
104
Demonstration, TWE, p. 93

These strategies are useful in a block scheduling format.

86D

CHAPTER

CHAPTER

4

Tying to Previous
Knowledge

The Structure of the Atom
What You’ll Learn
▲

Have students review the following concepts before studying this chapter. Chapter 1: matter, scientific laws, theories Chapter 3: elements, laws of definite and multiple proportions

Using the Photo

This CD-ROM is an editable
Microsoft® PowerPoint® presentation that includes:
• Section presentations
• Section checks
• Image bank
• Links to Chemistry Online
• All transparencies
• Animations
• Audio

Resource
Manager
Study Guide for Content
Mastery, p. 19 L2
Solving Problems: A Chemistry
Handbook, Section 4.1 L2
Section Focus Transparency 13 and Master L1 ELL

86

▲

scanning tunneling microscope in
1981 has made these images possible. ▲

Textbooks from 30 years ago claimed that no one could see atoms due to their extremely small size. Today, images of atoms are commonly seen in textbooks.
Why? The development of the

4

You will identify the experiments that led to the development of the nuclear model of atomic structure.
You will describe the structure of the atom and differentiate among the subatomic particles that comprise it.
You will explain the relationship between nuclear stability and radioactivity.

Why It’s Important
The world you know is made of matter, and all matter is composed of atoms.
Understanding the structure of the atom is fundamental to understanding why matter behaves the way it does.

Visit the Chemistry Web site at chemistrymc.com to find links about the structure of the atom.

Not only can individual atoms be seen, but scientists now have the ability to arrange them into patterns and simple devices. The atoms shown here have been arranged to form the Japanese kanji characters for atom.

86

Chapter 4

DISCOVERY LAB
Purpose
Students will observe the behavior of charged objects.

Teaching Strategies
• Any brand of plastic tape should give the same results.
• When rubbed, plastic and hard rubber gain a negative charge; glass, fur, and wool gain a positive charge.

• Remind students that charging the objects is a separation of existing electrical charge, not the creation of new electrical charge.
• Make a list of materials with the same charge as the tape stuck to the table and a list of those materials with the same charge as the pieces of tape that were stuck together.

Section 4.1

DISCOVERY LAB
Observing Electrical Charge

1 Focus

lectrical charge plays an important role in atomic structure and throughout chemistry. How can you observe the behavior of electrical charge using common objects?

E

Focus Transparency

Procedure

Before presenting the lesson, display
Section Focus Transparency 13 on the overhead projector. Have students answer the accompanying questions using Section Focus
Transparency Master 13. L1

1. Cut out small round pieces of paper using the hole punch and spread

them out on a table. Run a plastic comb through your hair. Bring the comb close to the pieces of paper. Record your observations.
2. Fold a 1-cm long portion of each piece of tape back on itself to

Materials metric ruler plastic comb hole punch paper 10-cm long piece of clear plastic tape (4)

form a handle. Stick two pieces of tape firmly to your desktop.
Quickly pull both pieces of tape off of the desktop and bring them close together so that their non-sticky sides face each other. Record your observations.

ELL

3. Firmly stick one of the remaining pieces of tape to your desktop.

Firmly stick the last piece of tape on top of the first. Quickly pull the pieces of tape as one from the desktop and then pull them apart. Bring the two tape pieces close together so that their nonsticky sides face each other. Record your observations.

Section
Focus

Transpare ncy 13

Gold Ato ms Use with

Chapter

4, Sectio

Analysis

Section

4.1

Objectives
• Compare and contrast the atomic models of
Democritus and Dalton.
• Define an atom.

Vocabulary
Dalton’s atomic theory atom Copyright
© Glencoe/Mc
Graw-Hill,
a division of the McGr aw-Hill Comp anies, Inc.

Use your knowledge of electrical charge to explain your observations.
Which charges are similar? Which are different? How do you know?

Early Theories of Matter
Perhaps you have never seen a photo of individual atoms as shown on the previous page, but chances are you’ve heard of atoms ever since you were in elementary school. From atom smashers and atomic power to the reality of the atomic bomb, you are already familiar with many modern atom-based processes. Surprisingly, the idea that matter is composed of tiny particles
(which we now call atoms) did not even exist a few thousand years ago. In fact, for more than a thousand years, great thinkers of their day argued against the idea that atoms existed. As you will see, the development of the concept of the atom and our understanding of atomic structure are fascinating stories involving scores of great thinkers and scientists.

The Philosophers
Science as we know it today did not exist several thousand years ago. No one knew what a controlled experiment was, and there were few tools for scientific exploration. In this setting, the power of mind and intellectual thought were considered the primary avenues to the truth. Curiosity sparked the interest of scholarly thinkers known as philosophers who considered the many mysteries of life. As they speculated about the nature of matter, many of the philosophers formulated explanations based on their own life experiences.
4.1 Early Theories of Matter

Expected Results

Analysis

Step 1: The pieces of paper are attracted to the comb.
Step 2: The pieces of tape repel each other as they are brought together.
Step 3: The pieces of tape attract each other as they are brought together.

Like electrical charges repel each other, whereas opposite electrical charges attract each other. The objects that repelled each other had similar charges. The objects that attracted each other had opposite charges.

87

1

2

Does the gold du st have bars? Ex the sam plain. e propert ies as th
Is there e gold a point at which so small gold can that the be particles gold? do not ret broken into parti cles ain the properties of

Chemistry:

Matter and
Change

Section

Focus Tran sparenc ies

Chapter Themes
The following themes from the
National Science Education
Standards are covered in this chapter. Refer to page 4T of the
Teacher Guide for an explanation of the correlations.
Systems, order, and organization
(UCP. 1); Evidence, models, and explanation (UCP.2); Change, constancy, and measurement
(UCP. 3); Evolution and equilibrium (UCP. 4); Form and function
(UCP. 5)

National Science Content Standards

Pages 86–87
UCP.1, UCP.2; A.1; B.1, B.2, B.4;
G.2, G.3

87

n 4.1

Figure 4-1

Quick Demo
Build a “mousetrap” sort of device using a marble that descends down a series of shoots, turns wheels, rings bells, and makes various noises. Many of these are available in children’s board games.
Shield the device with a cardboard box. You may wish to make part of the marble’s descent visible through cutout windows. Ask the students to draw what they think is going on inside the box based on their observations and have them draw parallels to discovery of the parts of the atom. Many Greek philosophers thought matter was formed of air, earth, fire, and water. They also associated properties with each of the four basic components of matter. The pairings of opposite properties, such as hot and cold, and wet and dry, mirrored the symmetry and balance the philosophers observed in nature. These early nonscientific and incorrect beliefs were not completely dispelled until the 1800s.

Air
Fire

Water
Earth

Dr

y

Co

ld

It wasn’t surprising then, that many of them concluded that matter was composed of things such as earth, water, air, and fire. See Figure 4-1. It was also commonly accepted that matter could be endlessly divided into smaller and smaller pieces. While these early ideas were creative, there was no method for testing their validity.
The Greek philosopher Democritus (460–370 B.C.) was the first person to propose the idea that matter was not infinitely divisible. He believed matter was made up of tiny individual particles called atomos, from which the
English word atom is derived. Democritus believed that atoms could not be created, destroyed, or further divided. Democritus and a summary of his ideas are shown in Figure 4-2.
While a fair amount of Democritus’s ideas do not agree with modern atomic theory, his belief in the existence of atoms was amazingly ahead of his time. Despite this, his ideas did not turn out to be a major step toward our current understanding of matter. Over time, Democritus’s ideas were met with criticism from other philosophers. “What holds the atoms together?” they asked. Democritus could not answer the question. Other criticisms came from
Aristotle (384–322 B.C.), one of the most influential Greek philosophers.
Aristotle is shown in Figure 4-3. He rejected the atomic “theory” entirely

Enrichment
Intrapersonal Prior to the

seventeenth century, scientists were known as natural philosophers who relied on the method of deductive reasoning to explain the world around them. In this method, philosophers observed events and extrapolated those observations through the use of reason back to other specific situations. The process was deductive because it went from general to specific, from universal to individual.
Sir Francis Bacon (1561–1626), an English lawyer and philosopher, argued passionately that although deductive reasoning might work in the realm of mathematics, the laws of science had to be induced, to be established as generalizations drawn from a vast amount of experimental observations. Have students compare inductive and deductive reasoning, citing examples of each. L2

W et t
Ho

Democritus’s Ideas
• Matter is composed of empty space through which atoms move.
• Atoms are solid, homogeneous, indestructible, and indivisible.
• Different kinds of atoms have different sizes and shapes.
• The differing properties of matter are due to the size, shape, and movement of atoms.
• Apparent changes in matter result from changes in the groupings of atoms and not from changes in the atoms themselves.

Figure 4-2
The Greek philosopher
Democritus (460–370 B.C.) proposed the concept of the atom more than two thousand years ago. 88

Chapter 4 The Structure of the Atom

CHEMISTRY JOURNAL
The Point of the Matter
Linguistic Show students a reproduction of any of Georges Seurat’s artwork which use the pointillistic technique. Sunday
Afternoon on the Grande Jatte is a good example. Show them the work at a distance

88

and then gradually decrease the viewing distance. Ask them to write how the artwork parallels Democritus’s ideas about the composition of matter. L2

Figure 4-3

Aristotle
• One of the most influential philosophers.
• Wrote extensively on many subjects, including politics, ethics, nature, physics, and astronomy. • Most of his writings have been lost through the ages.

The Greek philosopher Aristotle
(384–322 B.C.) was influential in the rejection of the concept of the atom.

because it did not agree with his own ideas on nature. One of Aristotle’s major criticisms concerned the idea that atoms moved through empty space. He did not believe that the “nothingness” of empty space could exist. Unable to answer the challenges to his ideas, Democritus’s atomic theory was eventually rejected.
In fairness to Democritus, it was impossible for him or anyone else of his time to determine what held the atoms together. More than two thousand years would pass before the answer was known. However, it is important to realize that Democritus’s ideas were just that—ideas and not science. Without the benefit of being able to conduct controlled experiments, Democritus could not test to see if his ideas were valid.
Unfortunately for the advancement of science, Aristotle was able to gain wide acceptance for his ideas on nature—ideas that denied the existence of atoms. Incredibly, the influence of Aristotle was so great and the development of science so primitive that his denial of the existence of atoms went largely unchallenged for two thousand years!

John Dalton
Although the concept of the atom was revived in the 18th century, it took the passing of another hundred years before significant progress was made. The work done in the 19th century by John Dalton (1766–1844), a schoolteacher in England, marks the beginning of the development of modern atomic theory. Dalton revived and revised Democritus’s ideas based upon the results of scientific research he conducted. The main points of Dalton’s atomic theory are shown in Figure 4-4.

Dalton’s Atomic Theory
• All matter is composed of extremely small particles called atoms.
• All atoms of a given element are identical, having the same size, mass, and chemical properties. Atoms of a specific element are different from those of any other element.
• Atoms cannot be created, divided into smaller particles, or destroyed.
• Different atoms combine in simple wholenumber ratios to form compounds.
• In a chemical reaction, atoms are separated, combined, or rearranged.

Extension
Intrapersonal John Dalton, a

practicing Quaker, left school at age 11 and then returned to teach at the ripe old age of 12! Dalton’s first love was meteorology, and he kept careful daily recordings of the weather for 57 years. His book
Meteorological Observations and
Essays (1793) qualified him as one of the pioneers in this area of study.
Such meteorological studies led
Dalton to think about the composition and properties of air. By studying the works of Democritus,
Boyle, and Proust, Dalton went on to formalize his own atomic theory.
This theory extended beyond simply gases—it encompassed all matter.
Dalton advanced his theory in his book New System of Chemical
Philosophy (1808). Have students research Dalton’s original writings on his atomic theory, citing examples of where the theory has been found to be in error and where it is still held to be true. L3

Enrichment
Obtain a video or videodisc of Ray and Charles Eames’ “Powers of
Ten,” a short and excellent dramatization that takes students on a journey from the subatomic level to intergalactic space. The video also provides an excellent review of scientific notation.

Assessment
Skill Given the number of
Figure 4-4
John Dalton’s (1766–1844) atomic theory was a breakthrough in our understanding of matter.

4.1 Early Theories of Matter

89

atoms in a penny, have students calculate the approximate volume of a single atom of copper. Then have them determine the number of atoms it would take to fill up any volume of interest, such as a 2-liter soda bottle. L2

DIFFERENTIATED INSTRUCTION
Hearing Impaired
CD-ROM
Chemistry: Matter and Change
Video: History of the Atomic
Theory

Visual-Spatial John Dalton believed that chemical formulas should be encoded as pictographs, in which each element had a particular drawing associated with it. Have students research Dalton’s pictographs and draw both the pictograph and its letter symbol counterpart. L2

National Science Content Standards

Pages 88–89
UCP.1, UCP.2; B.2; G.2, G.3

89

In-Text Question
Page 90 What similarities and differences can you find between the two theories? Similarities: matter composed of atoms; atoms are indestructible and indivisible; changes in matter are due to changes in groupings of atoms. Differences: Democritus states that matter is composed of empty space through which atoms move, whereas Dalton makes no such claims.

Figure 4-5
Dalton’s atomic theory explains the conservation of mass when a compound forms from its component elements. Atoms of elements A and B combine in a simple whole-number ratio, in this case two B atoms for each A atom, to form a compound.
Because the atoms are only rearranged in the chemical reactions, their masses are conserved.

Quick Demo
Place a 1-cm grid on an overhead projector. Place a clear lid or tray filled with water on top of the grid. Make an oleic acid solution (0.50% dissolved in methanol). Using a pipette and a small graduated cylinder, have students calculate the volume of a single drop of 0.50% oleic acid. Next, sprinkle the surface of the water with lycopodium or talcum powder. Place a single drop of the oleic acid in the middle of the water. The acid will spread out to the thickness of one molecule. Have students calculate the area by counting the cm squares. The density of oleic acid is 0.895 g/cm3 and the molecule has a length and width of 1/10 its thickness. Have students calculate the mass and thickness of the molecule.
Dispose of the oleic acid, lycopodium powder, talcum powder, and methanol by rinsing them down the drain with a large volume of water.

History
CONNECTION

A

lchemy, a popular pursuit during the Middle Ages, was the search for a way to transform common metals into gold.
“Scientists” who studied alchemy did not succeed in producing gold, but their experiments revealed the properties of various metals and helped advance our understanding of matter. During the seventeenth century some alchemists began focusing on identifying new compounds and reactions—these alchemists were probably the first true chemists.
By the twentieth century, scientists could transform atoms of some elements into atoms of other elements by bombarding them with neutrons. Even this process, however, has not succeeded in producing gold.

Tell students mass ϭ (volume)
(density) and thickness ϭ
(volume/area). L2
90

Pages 90–91
UCP.1, UCP.2; B.2; G.2, G.3

90

Atoms of element A
Total mass = mass A

Atoms of element B
Total mass = mass B

Compound composed of elements A and B
Total mass = (mass A + mass B)

The advancements in science since Democritus’s day served Dalton well, as he was able to perform experiments that allowed him to refine and verify his theories. Dalton studied numerous chemical reactions, making careful observations and measurements along the way. He was able to accurately determine the mass ratios of the elements involved in the reactions. Based on this research, he proposed his atomic theory in 1803. In many ways Democritus’s and Dalton’s theories are similar. What similarities and differences can you find between the two theories?
Recall from Chapter 3 that the law of conservation of mass states that mass is conserved in any process, such as a chemical reaction. Dalton’s atomic theory easily explains the conservation of mass in chemical reactions as being the result of the separation, combination, or rearrangement of atoms—atoms that are not created, destroyed, or divided in the process. The formation of a compound from the combining of elements and the conservation of mass during the process are shown in Figure 4-5. Dalton’s convincing experimental evidence and clear explanation of the composition of compounds and conservation of mass led to the general acceptance of his atomic theory.
Was Dalton’s atomic theory a huge step toward our current atomic model of matter? Yes. Was all of Dalton’s theory accurate? No. As is often the case in science, Dalton’s theory had to be revised as additional information was learned that could not be explained by the theory. As you will soon learn,
Dalton was wrong about atoms being indivisible (they are divisible into several subatomic particles) and about all atoms of a given element having identical properties (atoms of an element may have slightly different masses).

Defining the Atom
Many experiments since Dalton’s time have proven that atoms do actually exist. So what exactly then is the definition of an atom? To answer this question, consider a gold ring. Suppose you decide to grind the ring down into a pile of gold dust. Each fragment of gold dust still retains all of the properties of gold. If it were possible—which it is not without special equipment—you could continue to divide the gold dust particles into still smaller particles.
Eventually you would encounter a particle that could not be divided any further and still retain the properties of gold. This smallest particle of an element that retains the properties of the element is called an atom.
Just how small is a typical atom? To get some idea of its size, consider the population of the world. In the year 2000, the world population was approximately 6 000 000 000 (six billion) people. By comparison, a typical solid copper penny contains almost five billion times as many atoms of copper!
World population
6 000 000 000
Atoms in a penny
29 000 000 000 000 000 000 000

Chapter 4 The Structure of the Atom

Section 4.1

National Science Content Standards

ϩ

Assessment

1. Democritus could not explain what held

3. Dalton was wrong about atoms being

atoms together or verify that atoms moved through empty space.
2. An atom is the fundamental building block of nature. It is the smallest component of an element that exhibits all of the characteristic properties of that element.

indivisible (they are made up of subatomic protons, neutrons, and electrons) and about all atoms of a given element having identical properties (the masses of isotopes differ).
4. Democritus developed his ideas about atoms through intellectual thought, whereas Dalton developed his ideas by

3 Assess
Check for Understanding

a

Have students look through newspaper and magazine articles for examples of opinions, laws, and theories in science, government, economics, or any other area of interest. Have them copy the articles and highlight the examples. L2

b

The diameter of a single copper atom is 1.28 ϫ 10Ϫ10m. Placing six billion copper atoms (equal in number to the world’s population) side by side would result in a line of copper atoms less than one meter long.
You might think that because atoms are so small there would be no way to actually see them. However, an instrument called the scanning tunneling microscope allows individual atoms to be seen. Do the problem-solving
LAB on page 96 to analyze scanning tunneling microscope images and gain a better understanding of atomic size. As Figures 4-6a and 4-6b illustrate, not only can individual atoms be seen, scientists are now able to move individual atoms around to form shapes, patterns, and even simple machines. This capability has led to the exciting new field of nanotechnology. The promise of nanotechnology is molecular manufacturing—the atom-by-atom building of machines the size of molecules. As you’ll learn in later chapters, a molecule is a group of atoms that are bonded together and act as a unit. While this technology is not yet feasible for the production of consumer products, progress toward that goal has been made. To learn more about nanotechnology, read the Chemistry and Society at the end of this chapter.
The acceptance of atomic theory was only the beginning of our understanding of matter. Once scientists were fairly convinced of the existence of atoms, the next set of questions to be answered emerged. What is an atom like? How are atoms shaped? Is the composition of an atom uniform throughout, or is it composed of still smaller particles? While many scientists researched the atom in the 1800s, it was not until almost 1900 that answers to some of these questions were found. The next section explores the discovery of subatomic particles and the further evolution of atomic theory.

Section

4.1

Reteach

Figure 4-6 a This colorized scanning electron micrograph shows a microgear mechanism. b A mound of gold atoms (yellow, red, and brown) is easily discerned from the graphite substrate (green) it rests on.

Linguistic Read sections from the childhood favorite
The Cat in the Hat Comes Back, especially highlighting pages
32–63. Ask the students to write how the story parallels the development of the various atomic models discussed in this chapter. L2

Extension
Have students research how scanning tunneling microscopes (STM) and scanning probe microscopes
(SPM) function and why the technology is important. L2

CHEMLAB
ChemLab 4, located at the end of the chapter, can be used at this point in the lesson.

Assessment

Assessment

Performance Ask students
1.

Why were Democritus’s ideas rejected by other philosophers of his time?

2.

Define an atom using your own words.

3.

Which statements in Dalton’s original atomic theory are now considered to be incorrect? Describe how modern atomic theory differs from these statements.

chemistrymc.com/self_check_quiz

4.

Thinking Critically Democritus and Dalton both proposed the concept of atoms. Describe the method each of them used to reach the conclusion that atoms existed. How did Democritus’s method hamper the acceptance of his ideas?

5.

Comparing and Contrasting Compare and contrast the atomic theories proposed by Democritus and John Dalton.
4.1 Early Theories of Matter

91

to build a device such as the one mentioned in the Quick Demo on page 88. The contraptions can be shared in class where students can draw or describe how they think the contraptions work. Emphasize the connection between observation of a phenomenon and the model developed to explain it. Use the
Performance Task Assessment List for Invention in PASC, p. 45. L2
ELL

performing experiments and making careful measurements.
Because Democritus used only thought and not controlled experiments, he had no observable facts to support his ideas and no means of testing his theories.
5. Both believed: matter composed of extremely small particles called

atoms; all atoms of a given element are identical, but differ from the atoms of other elements; atoms could not be created, divided, or destroyed; apparent changes in matter result from changes in the groupings of atoms. Democritus further believed that matter is composed of empty space through which atoms move,

that different kinds of atoms come in different sizes and shapes, and that the differing properties of atoms are due to the size, shape, and movement of the atoms. Dalton further specified that different atoms combine in simple whole number ratios to form compounds.

91

Section 4.2

Section

4.2

Subatomic Particles and the
Nuclear Atom

1 Focus
Objectives

Focus Transparency

• Distinguish between the subatomic particles in terms of relative charge and mass.

Before presenting the lesson, display
Section Focus Transparency 14 on the overhead projector. Have students answer the accompanying questions using Section Focus
Transparency Master 14. L1

• Describe the structure of the nuclear atom, including the locations of the subatomic particles.

Vocabulary

ELL

14

an Atom
Models of with Chapter 4, Section 4.2
Use

anies, aw-Hill Comp of the McGr a division
Graw-Hill,
© Glencoe/Mc
Copyright

Inc.

nsparency
Focus Tra
Section

cathode ray electron nucleus proton neutron

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Focus Tran

Change
Matter and

2 Teach
Quick Demo
Use an electroscope or pith balls to show the transfer of electrons and the consequences. Rub a Lucite wand with fur or silk (this pulls off the electrons from the material) and then touch the wand to the electroscope or pith balls. The foil leaves of the electroscope
(or the pith balls themselves) will acquire a negative charge that will cause them to separate from each other. Reinforce that like charges repel.

ies

In 1839, American inventor Charles Goodyear accidentally heated a mixture of natural rubber and sulfur. The resulting reaction greatly strengthened the rubber. This new rubber compound revolutionized the rubber industry and was eventually used in the manufacturing of automobile tires. Accidental discoveries such as this have occurred throughout the history of science. Such is the case with the discovery of subatomic particles, the particles that make up atoms.

Discovering the Electron
Has your hair ever clung to your comb? Have you ever received a shock from a metal doorknob after walking across a carpeted floor? Observations such as these led scientists in the 1800s to look for some sort of relationship between matter and electric charge. To explore the connection, some scientists wondered how electricity might behave in the absence of matter. With the help of the recently invented vacuum pump, they passed electricity through glass tubes from which most of the air (and most of the matter) had been removed.
A typical tube used by researchers studying the relationship between mass and charge is illustrated in Figure 4-7. Note that metal electrodes are located on opposite ends of the tube. The electrode connected to the negative terminal of the battery is called the cathode, and the electrode connected to the positive terminal is called the anode.
One day while working in a darkened laboratory, English physicist Sir
William Crookes noticed a flash of light within one of the tubes. The flash was produced by some form of radiation striking a light-producing coating that had been applied to the end of the tube. Further work showed there were rays (radiation) traveling from the cathode to the anode within the tube. Because the ray of radiation originated from the cathode end of the tube, it became known as a cathode ray. The accidental discovery of the cathode ray led to the invention of one of the most important technological and social developments of the 20th century—the television. Television and computer monitor images are formed as radiation from the cathode strikes light-producing chemicals that coat the backside of the screen.

Figure 4-7
Examine the parts of a typical cathode ray tube. Note that the electrodes take on the charge of the battery terminal to which they are connected. The cathode ray travels from the cathode to the anode.

Tube filled with low pressure gas

Pages 92–93
UCP.1, UCP.2; A.2; B.1, B.4; G.1,
G.3

92

ϩ

Ϫ
Cathode
(Ϫ metal disk)

Anode
(ϩ metal disk)
Ϫ

Opening connected to a vacuum pump

ϩ

Voltage source 92

Chapter 4 The Structure of the Atom

Demonstration
Thomson’s Experiment
Purpose

National Science Content Standards

Cathode ray

To observe the characteristics of a cathode-ray tube
Materials

Crookes’s tube (cathode-ray tube); highvoltage DC power supply; bar magnet

Safety Precautions
Disposal Items can be saved and reused.
Procedure

Connect the two electrodes of the power supply to the ends of the tube using wires with alligator clips on each end. Turn on the power supply. CAUTION: The high-voltage power supply can cause severe electric shocks.
If the tube fails to light, increase the voltage from the power supply. If the tube still fails to light, turn off the power supply

2
Voltage source

Ϫ
Ϫ

ϩ

ϩ
S

Hole

The fact that the cathode ray is deflected in a magnetic field indicates that it is composed of charged particles. Enrichment
Linguistic The discovery of the inert gases in the 1890s led
Dimitri Mendeleev to speculate about ether, which scientists of the time believed was the component of empty space through which light traveled. Because Mendeleev predicted the existence and properties of elements heavier than hydrogen, he felt that he could also predict the existence of an element lighter than hydrogen (element X).
In his writings, he noted that element X was “capable of moving freely everywhere throughout the universe, with an atomic weight nearly one millionth that of hydrogen and traveling with a velocity of about 2250 kilometers per second.” He placed element X at the top of the family of inert gases.
Mendeleev resisted the concept of the electron because he felt that atoms had no further internal structure.
Have students investigate and write about other scientists who, while credited with one or more discoveries, were in error with some of their other theories. L3

N
Magnet

Anode
Cathode

1

Gas at low pressure

A tiny hole located in the center of the anode produces a thin beam of electrons. A phosphor coating allows the position of the beam to be determined as it strikes the end of the tube.
Because altering the gas in the tube and the material used for the cathode have no effect on the cathode ray, the particles in the ray must be part of all matter. 3
Electrically
charged plates

Ϫ

ϩ

ϩ

Because the cathode ray is deflected toward the positively charged plate by an electric field, the particles in the ray must have a negative charge.

Ϫ

Figure 4-8

Scientists continued their research using cathode ray tubes, and by the end of the 1800s they were fairly convinced of the following:
• Cathode rays were actually a stream of charged particles.
• The particles carried a negative charge. (The exact value of the negative charge was not known, however.)

Multiple experiments helped determine the properties of cathode rays.

Because changing the type of electrode or varying the gas (at very low pressure) in the cathode ray tube did not affect the cathode ray produced, it was concluded that the ray’s negative particles were found in all forms of matter.
These negatively charged particles that are part of all forms of matter are now called electrons. The range of experiments used to determine the properties of the cathode ray are shown in Figure 4-8.
In spite of the progress made from all of the cathode ray tube experiments, no one had succeeded in determining the mass of a single cathode ray particle. Unable to measure the particle’s mass directly, English physicist
J.J. Thomson (1856–1940) began a series of cathode ray tube experiments in the late 1890s to determine the ratio of its charge to its mass. By carefully measuring the effect of both magnetic and electric fields on a cathode ray,
Thomson was able to determine the charge-to-mass ratio of the charged particle. He then compared that ratio to other known ratios. Thomson concluded that the mass of the charged particle was much less than that of a hydrogen atom, the lightest known atom. The conclusion was shocking because it meant there were particles smaller than the atom. In other words,
Dalton was wrong: Atoms were divisible into smaller subatomic particles.
Because Dalton’s atomic theory had become so widely accepted, and because
Thomson’s conclusion was so revolutionary, many fellow scientists found it hard to believe this new discovery. But Thomson was correct. He had identified the first subatomic particle—the electron.
The next significant development came in 1909, when an American physicist named Robert Millikan (1868–1953) determined the charge of an electron.
4.2 Subatomic Particles and the Nuclear Atom

and reverse the wire connections on the tube. Have students observe the electron beam on the fluorescent screen. Deflect the beam with the magnet. Reverse the magnet’s field by flipping it around and use it to show deflection of the beam in the opposite direction.
Results

Students will see evidence of an electron beam. The phosphors on the plate glow

along the path of the beam. Point out that
Thomson observed deflection by means of both magnets and charged plates.
Analysis

1. What properties account for the bending of the cathode ray in a magnetic field? The ray has an electrical charge.

2. What is the function of the electromagnets on a TV picture tube? They cause

Resource
Manager
Study Guide for Content
Mastery, p. 20 L2
Solving Problems: A Chemistry
Handbook, Section 4.2 L2
Section Focus Transparency 14 and Master L1 ELL
Teaching Transparency 11 and
Master L2 ELL

93

the electron beam to move back and forth, forming an image on the inside face of the luminescent chemicalcoated screen.

Assessment
Skill Have students research an article describing how a CRT or television works. L2
93

Quick Demo
Show the students a CRT with a paddlewheel or use a photograph of one (lab equipment catalogs are a good source of such photographs). Emphasize to the students that the only way for the paddlewheel to move uphill against gravity is to be bombarded by a stream of particles emanating from the cathode.

Applying Chemistry

Matter containing evenly distributed positive
Electrons
charge

Ϫ

Ϫ
Ϫ

Ϫ

Ϫ

Ϫ

Ϫ

Ϫ
Ϫ

Ϫ

Ϫ

Ϫ

Ϫ
Figure 4-9

J.J. Thomson’s plum pudding atomic model proposed that negatively charged electrons were distributed throughout a uniform positive charge.

A more sophisticated version of the cathode ray tube can be found in most homes in the form of a television set or computer monitor. A television uses magnets to direct the electron beam back and forth across the backside of its screen. The screen is coated with luminescent chemicals that produce the colors we see when it is illuminated by the electron beam.

Mass of an electron ϭ 9.1 ϫ 10Ϫ28 g ϭ ᎏ1ᎏ mass of a hydrogen atom
1840
As you can see, the mass of an electron is extremely small.
The existence of the electron and the knowledge of some of its properties raised some interesting new questions about the nature of atoms. It was known that matter is neutral. You know matter is neutral from everyday experience; you do not receive an electrical shock (except under certain conditions) when you touch an object. If electrons are part of all matter and they possess a negative charge, how is it that all matter is neutral? Also, if the mass of an electron is so extremely small, what accounts for the rest of the mass in a typical atom?
In an attempt to answer these questions, J.J. Thomson proposed a model of the atom that became known as the plum pudding model. As you can see in Figure 4-9, Thomson’s model consisted of a spherically shaped atom composed of a uniformly distributed positive charge within which the individual negatively charged electrons resided. A more modern name for this model might be the chocolate-chip cookie dough model, where the chocolate chips are the electrons and the dough is the uniformly distributed positive charge.
As you are about to learn, the plum pudding model of the atom did not last for very long.

The Nuclear Atom

Figure 4-10

CD-ROM
Chemistry: Matter and Change
Experiment: Discovery of the Electron
Demonstration:
Thomson’s Experiment
Animation:
Rutherford’s Gold

So good was Millikan’s experimental setup and technique that the charge he measured almost one hundred years ago is within 1% of the currently accepted value. This charge has since been equated to a single unit of negative charge; in other words, a single electron carries a charge of 1–. Knowing the electron’s charge and using the known charge-to-mass ratio, Millikan calculated the mass of a single electron.

Ernest Rutherford expected most of the fast-moving and relatively massive alpha particles to pass straight through the gold atoms.
He also expected a few of the alpha particles to be slightly deflected by the electrons in the gold atoms.
Evenly distributed positive charge

Electrons

Ϫ

Ϫ

Alpha particle path

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Ϫ
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Ϫ

94

Ϫ
Ϫ
Ϫ

The story of the atom continues with the role played by Ernest Rutherford
(1871–1937). As a youth, Rutherford, who was born in New Zealand, placed second in a scholarship competition to attend the prestigious Cambridge
University in England. He received a fortunate break when the winner of the competition decided not to attend. By 1908, Rutherford won the Nobel Prize in chemistry and had many significant discoveries to his credit.
In 1911 Rutherford became interested in studying how positively charged alpha particles (radioactive particles you will learn more about later in this chapter) interacted with solid matter. A small group of scientists that included
Rutherford designed and conducted an experiment to see if alpha particles would be deflected as they passed through a thin foil of gold. In the experiment, a narrow beam of alpha particles was aimed at a thin sheet of gold foil.
A zinc sulfide coated screen surrounding the gold foil produced a flash of light whenever it was struck by an alpha particle. By noting where the flashes occurred, the scientists could determine if the atoms in the gold foil deflected the alpha particles.
Rutherford was aware of Thomson’s plum pudding model of the atom and expected only minor deflections of the alpha particles. He thought the paths of the massive (relative to electrons) and fast-moving alpha particles would be only slightly altered by a nearby encounter or collision with an electron.
And because the positive charge within the gold atoms was thought to be uniformly distributed, he thought it would not alter the paths of the alpha particles either. Figure 4-10 shows the results Rutherford anticipated from the experiment. After a few days of testing, Rutherford and his fellow scientists were amazed to discover that a few of the alpha particles were deflected at

Chapter 4 The Structure of the Atom

CHEMISTRY JOURNAL

National Science Content Standards

Pages 94–95
UCP.1, UCP.2; A.2; B.1, B.4; G.2, G.3

94

DIFFERENTIATED INSTRUCTION

Comparing Models

Gifted

Visual-Spatial Have students use their journals to sketch the structure of an atom according to the plum pudding and nuclear atomic models. Have them label the subatomic particles in each sketch and list the key points associated with each model. They should clearly label that the nuclear model is the currently accepted model of the atom. L2 ELL

Linguistic Interested students may wish to read and report on the oil drop experiment performed by Robert
Millikan. L3

Figure 4-11

Lead block containing an alpha particle emitting source
Alpha particle deflected at a large angle

As Rutherford expected, most all of the alpha particles passed straight through the gold foil, without deflection. Surprisingly, however, some alpha particles were scattered at small angles, and on a few occasions they were deflected at very large angles. Beam of alpha particles

Assessment
Portfolio Ask students to

react to the following statement:
Most of matter is nothing. Have them use specific experimental evidence discussed in this section.
Students’ work can be placed in their portfolios. L2 P

Gold foil

Resource
Manager

Zinc sulfide coated screen
Most alpha particles pass through the foil with little or no deflection Alpha particle deflected at a small angle

Teaching Transparency 12 and
Master L2 ELL

very large angles. Several particles were even deflected straight back toward the source of the alpha particles. Rutherford likened the results to firing a large artillery shell at a sheet of paper and having the shell come back and hit you!
These results, shown in Figure 4-11, were truly astounding.
Rutherford concluded that the plum pudding model was incorrect because it could not explain the results of the gold foil experiment. He set out to develop a new atomic model based upon his findings. Considering the properties of the alpha particles and the electrons, and the frequency of the deflections, he calculated that an atom consisted mostly of empty space through which the electrons move. He also concluded that there was a tiny, dense region, which he called the nucleus, centrally located within the atom that contained all of an atom’s positive charge and virtually all of its mass. Because the nucleus occupies such a small space and contains most of an atom’s mass, it is incredibly dense. Just how dense? If a nucleus were the size of the dot in the exclamation point at the end of this sentence, its mass would be approximately as much as that of 70 automobiles!
According to Rutherford’s new nuclear atomic model, most of an atom consists of electrons moving rapidly through empty space. The electrons move through the available space surrounding the nucleus and are held within the atom by their attraction to the positively charged nucleus. The volume of space through which the electrons move is huge compared to the volume of the nucleus. A typical atom’s diameter, which is defined by the volume of space through which the electrons move, is approximately 10 000 times the diameter of the nucleus. To put this in perspective, if an atom had a diameter of two football fields, the nucleus would be the size of a nickel!
The concentrated positive charge in the nucleus explains the deflection of the alpha particles—the repulsive force produced between the positive nucleus and the positive alpha particles causes the deflections. Alpha particles closely approaching the nucleus were deflected at small angles, while alpha particles directly approaching the nucleus were deflected at very large angles. You can see in Figure 4-12 how Rutherford’s nuclear atomic model explained the

Figure 4-12
Rutherford’s nuclear model of the atom explains the results of the gold foil experiment. Most alpha particles pass straight through, being only slightly deflected by electrons, if at all. The strong force of repulsion between the positive nucleus and the positive alpha particles causes the large deflections.
Electrons

Ϫ

Ϫ

Ϫ

Alpha particle path

Ϫ

ϩ

Ϫ

Ϫ
Ϫ

Ϫ

Ϫ
Ϫ
Ϫ

Nucleus

Ϫ

4.2 Subatomic Particles and the Nuclear Atom

95

DIFFERENTIATED INSTRUCTION
Learning Disabled
Kinesthetic Consider creating a gold foil “shooting gallery” within a part of your classroom to conceptually recreate aspects of Rutherford’s experiment. Using string, hang tennis balls from the ceiling at
2-foot intervals (horizontally and vertically) so that they form a small rectangular field.
These tennis balls represent nuclei of the gold atoms in the foil. Have students take

turns throwing a tennis ball “alpha particle” through the field of tennis ball “gold atoms.” Record the number of throws that go through versus those that are deflected.
Ask students to compare the in-classroom setup and results with the those of
Rutherford’s gold foil experiment. L1 ELL

95

problem-solving LAB
Purpose

Students will develop a measurement scale conversion factor and use it to make measurements from an STM image. They will also make calculations based on those measurements. Topic: Atomic Structure
To learn more about the discovery of atomic structure, visit the Chemistry Web site at chemistrymc.com
Activity: Research Ernest
Rutherford’s work. Write a newspaper article announcing his model of the atom.

Process Skills

Measuring, interpreting scientific illustrations, using numbers, drawing conclusions
Teaching Strategies

results of the gold foil experiment. The nuclear model also explains the neutral nature of matter: the positive charge of the nucleus balancing the negative charge of the electrons. However, the model still could not account for all of the atom’s mass. Another 20 years would pass before this mystery was solved.

Completing the Atom—The Discovery of
Protons and Neutrons
By 1920, eight years after his revolutionary gold foil experiment, Rutherford had refined the concept of the nucleus. He concluded that the nucleus contained positively charged particles called protons. A proton is a subatomic particle carrying a charge equal to but opposite that of an electron; that is, a proton has a positive charge of 1+.
In 1932, Rutherford’s coworker, English physicist James Chadwick
(1891–1974), showed that the nucleus also contained another subatomic particle, a neutral particle called the neutron. A neutron has a mass nearly equal to that of a proton, but it carries no electrical charge. Thus, three subatomic particles are the fundamental building blocks from which all atoms are

• Ask students how a mileage scale works on a road map.

• Other STM images may be used as an introduction to the techniques of developing a scale from the image and applying geometric relationships to determine radii, diameters, etc.
Thinking Critically

1. The real dimensions of the image are 18.1 nm wide by
19.0 nm high. Measuring the edges of the image with a ruler gives the dimensions in centimeters. By dividing the centimeter measurement by the scale shown in the photo, a conversion factor can be determined. The distance between adjacent nuclei is about 0.7 nm.
2. The waves in the middle seem to indicate the presence of an electron. The distance between adjacent nuclei is about 93 nm.

Assessment
Performance Using

photos that have a feature with a commonly known or determinable length in them, have students determine size of other objects in the photos. Use the Performance Task
Assessment List for Making
Observations and Inferences in
PASC, p. 17. L2

problem-solving LAB
Interpreting STM Images
Measuring The invention of the scanning tunneling microscope (STM) in 1981 gave scientists the ability to visualize individual atoms, and also led to their being able to manipulate the positions of individual atoms. Use the information shown in the STM images to interpret sizes and make measurements.

Analysis
Figure A is an STM image of silicon atoms that have been bonded together in a hexagonal pattern. The image is of an area 18.1nm wide by
19.0 nm high (1 nm = 1 x 10 Ϫ9 m).
Figure B is an STM image of 48 iron atoms that have been arranged into a circular "corral."
The corral has a diameter of 1426 nm. There is a single electron trapped inside the "corral."

Figure A

Thinking Critically
1. Using a metric ruler and the dimensions of
Figure A given above, develop a scale for making measurements off of the image. Use your scale to estimate the distance between adjacent silicon nuclei forming a hexagon.
2. What evidence is there that an electron is trapped inside the “corral” of iron atoms in Figure B? Estimate the distance between adjacent iron atoms. (Hint: Use the number of atoms and the formula circumference ϭ ␲ ϫ diameter.)

Figure B
96

Chapter 4 The Structure of the Atom

Internet Address Book
Note Internet addresses that you find useful in the space below for quick reference.

National Science Content Standards

Pages 96–97
UCP.1, UCP.2; A.2; B.1; G.2, G.3

96

Table 4-1

Electron cloud

Relative electrical charge

Location

3 Assess

Nucleus

Properties of Subatomic Particles
Relative
mass

Actual mass (g)

Particle

Symbol

Electron

eϪ

In the space surrounding the nucleus

1Ϫ

1
ᎏᎏ
1840

9.11 ϫ 10Ϫ28

Proton

pϩ

In the nucleus

1ϩ

1

1.673 ϫ 10Ϫ24

Neutron

n0

In the nucleus

0

1

1.675 ϫ 10Ϫ24

Check for Understanding
Ask students to compare functions and dimensions of the basic unit of chemistry, the atom, with the basic unit of biology, the cell.

Reteach
Kinesthetic Have groups

made—the electron, the proton, and the neutron. Together, electrons, protons, and neutrons account for all of the mass of an atom. The properties of electrons, protons, and neutrons are summarized in Table 4-1.
You know an atom is an electrically neutral particle composed of electrons, protons, and neutrons. Atoms are spherically shaped, with a tiny, dense nucleus of positive charge surrounded by one or more negatively charged electrons. Most of an atom consists of fast-moving electrons traveling through the empty space surrounding the nucleus. The electrons are held within the atom by their attraction to the positively charged nucleus. The nucleus, which is composed of neutral neutrons (hydrogen’s single-proton nucleus is an exception) and positively charged protons, contains all of an atom’s positive charge and more than 99.97% of its mass. Since an atom is electrically neutral, the number of protons in the nucleus equals the number of electrons surrounding the nucleus. The features of a typical atom are shown in Figure 4-13.
To gain more perspective on the size of typical atoms, do the CHEMLAB at the end of this chapter.
Subatomic particle research is still a major interest of modern scientists.
In fact, the three subatomic particles you have just learned about have since been found to have their own structures. That is, they contain sub-subatomic particles. These particles will not be covered in this textbook because it is not understood if or how they affect chemical behavior. As you will learn in coming chapters, behavior can be explained by considering only an atom’s electrons, protons, and neutrons.
You should now have a solid understanding of the structure of a typical atom. But what makes an atom of one element different from an atom of another element? In the next section, you’ll find out.

Proton

Neutron

of students build the various atomic models using clay, construction paper, gum drops, or any other materials of their choice. L1 ELL

Extension
Visual-Spatial Have
Figure 4-13
Atoms consist of a “cloud” of fast moving, negatively charged electrons surrounding a tiny, extremely dense nucleus containing positively charged protons and neutral neutrons. The nucleus contains virtually all of the atom’s mass, but occupies only about one ten-thousandth the volume of the atom.

students make a timeline of the events leading to our current understanding of the atom. The timeline should include anyone involved in the evolution of our understanding of the atom and also include a brief description of any related experiments. L2

Assessment
Skill Using the mass of an

LAB
See page 953 in Appendix E for
Comparing Atom Sizes

electron as the basic unit of mass, have students set up a table that relates the masses of the other subatomic particles, other atoms, and other objects of their choosing.
L2

Section

4.2

Assessment
9. Thomson’s plum pudding

6.

Briefly evaluate the experiments that led to the conclusion that electrons were negatively charged particles found in all matter.

9.

7.

Describe the structure of a typical atom. Be sure to identify where each subatomic particle is located.

10.

8.

Make a table comparing the relative charge and mass of each of the subatomic particles.

chemistrymc.com/self_check_quiz

Section 4.2

Thinking Critically Compare and contrast
Thomson’s plum pudding atomic model with
Rutherford’s nuclear atomic model.
Graphing Make a timeline graph of the development of modern atomic theory. Be sure to include the discovery of each subatomic particle.

4.2 Subatomic Particles and the Nuclear Atom

97

Assessment

6. The deflection toward positively charged plates demonstrated the negatively charged nature of electrons; the fact that changing the type of electrode or the type of gas used in the cathode ray tube did not affect the ray produced led to the conclusion that electrons are present in all matter.

7. A typical atom consists of a central, small, dense nucleus containing protons and neutrons. The nucleus is surrounded by a cloud of negatively charged electrons.

8. Particle
Electron
Proton
Neutron

Relative charge

Relative mass

Ϫ1 ϩ1 0

1/1840
1
~1

model describes atoms as spherical particles with uniformly distributed positive charge in which individual, negatively charged, electrons are located. In contrast,
Rutherford’s model states that an atom is mostly empty
National Science Content Standards space, with a small, dense, central nucleus containing all of an atom’s positive charge and most of its mass. The negatively charged electrons move through the empty space and are held in the atom by their attraction to the positively charged nucleus. 10. Timelines must include
Democritus, Aristotle,
Dalton, Crookes, Thomson,
Millikan, and Rutherford.

97

Section

Section 4.3

4.3

Objectives

1 Focus

• Explain the role of atomic number in determining the identity of an atom.

Focus Transparency

• Define an isotope and explain why atomic masses are not whole numbers.

Before presenting the lesson, display
Section Focus Transparency 15 on the overhead projector. Have students answer the accompanying questions using Section Focus Transparency
Master 15. L1 ELL

• Calculate the number of electrons, protons, and neutrons in an atom given its mass number and atomic number. Vocabulary nsparency Focus Tra
Section

15

atomic number isotope mass number atomic mass unit (amu) atomic mass

ts d Elemen
Atoms an h Chapter 4, Section 4.3
Use wit

Sodium
11

Na

anies, aw-Hill Comp of the McGr a division
Graw-Hill,
© Glencoe/Mc
Copyright

Inc.

22.990

Sulfur
16

How Atoms Differ
Look at the periodic table on the inside back cover of this textbook. As you can see, there are more than 110 different elements. This means that there are more than 110 different kinds of atoms. What makes an atom of one element different from an atom of another element? You know that all atoms are made up of electrons, protons, and neutrons. Thus, you might suspect that atoms somehow differ in the number of these particles. If so, you are correct.

Atomic Number
Not long after Rutherford’s gold foil experiment, the English scientist Henry
Moseley (1887–1915) discovered that atoms of each element contain a unique positive charge in their nuclei. Thus, the number of protons in an atom identifies it as an atom of a particular element. The number of protons in an atom is referred to as the element’s atomic number. Look again at the periodic table and you will see that the atomic number determines the element’s position in the table. Consider hydrogen, located at the top left of the table. The information provided by the periodic table for hydrogen is shown in Figure 4-14.
Note that above the symbol for hydrogen (H), you see the number 1. This number, which corresponds to the number of protons in a hydrogen atom, is the atomic number of hydrogen. Hydrogen atoms always contain a single proton. Moving across the periodic table to the right, you’ll next come to helium
(He). Helium has an atomic number of 2, and thus has two protons in its nucleus. The next row begins with lithium (Li), atomic number 3, followed by beryllium (Be), atomic number 4, and so on. As you can see, the periodic table is organized left-to-right and top-to-bottom by increasing atomic number. How many protons does a gold atom contain? A silver atom?
Remember that because all atoms are neutral, the number of protons and electrons in an atom must be equal. Thus, once you know the atomic number of an element, you know both the number of protons and the number of electrons an atom of that element contains.

S

Atomic number ϭ number of protons ϭ number of electrons

32.066

? properties in their nts differ two eleme s differ? do these element 1 How ese two th of e atoms might th
2 How
Section

Chemistry:

sparenc
Focus Tran

For instance, an atom of lithium, atomic number of 3, contains three protons and three electrons. How many electrons does an atom of element 97 contain?

ies

Change
Matter and

2 Teach
In-Text Questions
Page 98 How many protons does a gold atom contain? 79 A silver atom? 47 How many electrons does an atom of element 97 contain? 97

Figure 4-14
The atomic number of an element equals the positive charge contained in its nucleus.

H
1.008

Resource
Manager
Study Guide for Content
Mastery, pp. 21–23 L2
Solving Problems: A Chemistry
Handbook, Section 4.3 L2
Section Focus Transparency 15 and Master L1 ELL

National Science Content Standards

Pages 98–99
UCP.1, UCP.2; B.1; B.2; G.2

98

Hydrogen
1

Chemical name
Atomic number
Chemical symbol
Average atomic mass

Hydrogen, with an atomic number of 1, is the first element in the periodic table. A hydrogen atom has one proton and a charge of 1ϩ in its nucleus.

98

Chapter 4 The Structure of the Atom

DIFFERENTIATED INSTRUCTION
Learning Disabled
Kinesthetic Use three different colors of gum drops to allow students to build models of atoms that show the relationship of protons, neutrons, and electrons.
Be sure to emphasize that the scaling of the

model is in no way appropriate; electrons are 1/1840 the size of protons and neutrons, and the diameter of an atom is about 10 000 times the size of its nucleus. L1 ELL

EXAMPLE PROBLEM 4–1

Enrichment

Using Atomic Number

Of the scientists working under
Rutherford at Cambridge University,
Henry Moseley was the youngest.
He beamed X rays at samples of the different elements to find that the wavelength of the resultant X rays decreased with increasing atomic weight. Moseley attributed this to the increasing number of electrons in the atom and, consequently, the increasing number of protons in the nucleus. This led to the final arrangement of Mendeleev’s periodic table—elements arranged by increasing number of protons
(atomic number). This arrangement solved problems with the sequence of elements with nearly similar atomic masses and clearly identified
“holes” in the table for yet-to-be discovered elements; as of 1914, there were seven undiscovered elements. Moseley’s method was used to refute Urbain’s claim of discovering a new element that he called “celtium” and to support
Hevesey’s discovery of hafnium.
Have students compare Mendeleev’s periodic table with Moseley’s periodic table and identify discrepancies in the sequence of elements. L3

Complete the following table.

Composition of Several Elements
Element

Atomic number

Protons

Electrons

a.

Pb

82

___

___

b.

___

___

8

___

c.

___

___

___

30

1. Analyze the Problem
You are given the information in the table. Apply the relationship among atomic number, number of protons, and number of electrons to complete most of the table. Once the atomic number is known, use the periodic table to identify the element.
2. Solve for the Unknown
Apply the atomic number relationship and then consult the periodic table to identify the element.
a. Atomic number ϭ number of protons ϭ number of electrons
82 ϭ number of protons ϭ number of electrons
Element 82 is lead (Pb).
b. Atomic number ϭ number of protons ϭ number of electrons
Atomic number ϭ 8 ϭ number of electrons
Element 8 is oxygen (O).
c. Atomic number ϭ number of protons ϭ number of electrons
Atomic number ϭ number of protons ϭ 30
Element 30 is zinc (Zn).
The completed table is shown below.

Composition of Several Elements
Element

Atomic number

Protons

Electrons

a.

Pb

82

82

82

b.

O

8

8

8

c.

Zn

30

30

30

3. Evaluate the Answer

PROBLEMS

The answers agree with atomic numbers and element symbols given in the periodic table.

PRACTICE PROBLEMS
11. How many protons and electrons are in each of the following atoms?
a. boron
b. radon

c. platinum
d. magnesium

12. An atom of an element contains 66 electrons. What element is it?

e!
Practic

For more practice with problems using atomic numbers, go to
Supplemental Practice
Problems in Appendix A.

13. An atom of an element contains 14 protons. What element is it?

4.3 How Atoms Differ

Have students refer to Appendix
D for complete solutions to
Practice Problems.
11. a. boron, 5
b. radon, 86
c. platinum, 78
d. magnesium, 12
12. dysprosium
13. silicon

99

A Woman with Designs
In 1948, a high school career aptitude test indicated Patsy Sherman was best suited to be a housewife. Instead, she decided to attend college, where she earned degrees in chemistry and math.
After graduation, she got a temporary job at the 3M Corporation in St. Paul,
Minnesota.

One day, a lab assistant accidentally spilled some chemicals on Patsy’s shoes.
She noticed that the site of the spill was waterproof and stain-resistant. She decided to pursue the chemical mixture, which later was marketed as the fabric protectant Scotchgard. The protectant consists of a rubbery molecule, one side of which is sticky, allowing it to cling to

fabric, while the other is slippery enough to repel stains.
3M liked the idea so well that it awarded Sherman with a full-time position. Scotchgard became one of 3M’s most profitable products. Sherman has gone on to earn 16 U.S. patents and serves on the board of the National
Inventors Hall of Fame.
99

Isotopes and Mass Number

Identifying
Misconceptions
Students may think that isotopes contain different numbers of electrons and protons.
Uncover the Misconception
Ask students if and how isotopes of an atom differ from each other.
Demonstrate the Concept
Explain to students that an atom’s identity is defined solely by the number of protons in its nucleus. The number of neutrons may vary, resulting in the existence of different isotopes. The number of electrons in the neutral isotopes is the same. Reinforce that various ions (charged particles) of an element may have different numbers of electrons.
Assess New Knowledge
Ask a student to show on the board how three isotopes of oxygen (oxygen-16, oxygen-17, and oxygen-18) are the same and how they are different. All

Mass number

107
47

Ag

109
47

Ag

Atomic number

Figure 4-15
Ag is the chemical symbol for the silver used in these coins.
The silver in each coin is comprised of 51.84% silver-107
(107
47 Ag) isotopes and 48.16% silver-109 (109
47 Ag) isotopes.

isotopes have 8 protons, but the three isotopes differ by having 8, 9, and 10 neutrons respectively. Earlier you learned that Dalton’s atomic theory was wrong about atoms being indivisible. It was also incorrect in stating that all atoms of a particular element are identical. While it is true that all atoms of a particular element have the same number of protons and electrons, the number of neutrons on their nuclei may differ. For example, there are three different types of potassium atoms. All three types contain 19 protons (and thus 19 electrons). However, one type of potassium atom contains 20 neutrons, another contains 21 neutrons, and still another 22 neutrons. Atoms such as these, with the same number of protons but different numbers of neutrons, are called isotopes.
In nature most elements are found as a mixture of isotopes. Usually, no matter where a sample of an element is obtained, the relative abundance of each isotope is constant. For example, in a banana, which is a rich source of potassium,
93.25% of the potassium atoms have 20 neutrons, 6.7302% will have 22 neutrons, and a scant 0.0117% will have 21 neutrons. In another banana, or in a totally different source of potassium, the percentage composition of the potassium isotopes will still be the same.
As you might expect, the isotopes do differ in mass. Isotopes containing more neutrons have a greater mass. In spite of differences in mass and the number of neutrons, isotopes of an atom have essentially the same chemical behavior.
Why? Because, as you’ll learn in greater detail later in this textbook, chemical behavior is determined by the number of electrons an atom has, not by its number of neutrons and protons. To make it easy to identify each of the various isotopes of an element, chemists add a number after the element’s name. The number that is added is called the mass number, and it represents the sum of the number of protons and neutrons in the nucleus. For example, the potassium isotope with 19 protons and 20 neutrons has a mass number of 39 (19 ϩ 20 ϭ
39), and the isotope is called potassium-39. The potassium isotope with 19 protons and 21 neutrons has a mass number of 40 (19 ϩ 21 ϭ 40), and is called potassium-40. What is the mass number and name of the potassium isotope with
19 protons and 22 neutrons?
Chemists often write out isotopes using a shortened type of notation involving the chemical symbol, atomic number, and mass number, as shown in
Figure 4-15. Note that the mass number is written as a superscript to the left of the chemical symbol, and the atomic number is written as a subscript to the left of the chemical symbol. The three potassium isotopes you have just learned

Figure 4-16

Figure Caption Questions
Figure 4-16 How do their masses compare? Their chemical properties? From K-39 to K-40 to K-41,

The three naturally occurring potassium isotopes are potassium-39, potassium-40, and potassium-41. How do their masses compare? Their chemical properties? their masses differ by the mass of a single neutron. The chemical properties of the isotopes are the same.

19pϩ
20n0

and the isotope is potassium-41.

100

Chapter 4 The Structure of the Atom

CHEMISTRY JOURNAL
Useful Isotopes

National Science Content Standards

Pages 100–101
UCP.1, UCP.2; B.2

100

Potassium-40

Potassium-41

19
20
19

19
21
19

19
22
19

19eϪ

In-Text Question
Page 100 What is the mass number and name of the potassium isotope with 19 protons and 22 neutrons? The mass number is 41

Potassium-39
Protons
Neutrons
Electrons

Linguistic Have students research the uses of isotopes in medicine and radiochemical dating. In their journals, have them describe how radiochemical dating makes use of radioactive isotopes.
Also, have them list at least three isotopes used in medicine and describe how those isotopes are used. L2

19eϪ

19pϩ
21n0

19eϪ

19pϩ
22n0

about are summarized in Figure 4-16. The number of neutrons in an isotope can be calculated from the atomic number and mass number.

Extension

Number of neutrons ϭ mass number Ϫ atomic number

EXAMPLE PROBLEM 4–2
Using Atomic Number and Mass Number

Isotope Composition Data

A chemistry laboratory has analyzed the composition of isotopes of several elements. The composition data is given in the table at the right.
Data for one of neon’s three isotopes is given in the table. Determine the number of protons, electrons, and neutrons in the isotope of neon. Name the isotope and give its symbol.

Element

Atomic number Mass number a. Neon

10

22

b. Calcium

20

46

1. Analyze the Problem

c. Oxygen

8

17

d. Iron

26

57

e. Zinc

30

64

f. Mercury

80

204

You are given some data for neon in the table. The symbol for neon can be found from the periodic table. From the atomic number, the number of protons and electrons in the isotope are known. The number of neutrons in the isotope can be found by subtracting the atomic number from the mass number.
Known

Unknown

Element: neon
Atomic number ϭ 10
Mass number ϭ 22

Number of protons, electrons, and neutrons ϭ ?

Familiarize your students with the
Table of Isotopes in the CRC
Handbook of Chemistry and
Physics. At this point, they should be able to identify how many isotopes are found per element and understand how the isotope notation is listed. Check student understanding by quizzing them on facts they look up in the table.

Assessment

Name of isotope ϭ ?
Symbol for isotope ϭ ?

2. Solve for the Unknown
The number of protons equals the number of electrons which equals the atomic number.

Performance Have students use the overhead projector and a set of colored dots (cut from colored acetate sheets) representing protons, neutrons, and electrons to build correct configurations for various isotopes. Use the
Performance Task Assessment List for Model in PASC, p. 51. L1
ELL

Number of protons ϭ number of electrons = atomic number ϭ 10
Use the atomic number and the mass number to calculate the number of neutrons.

PROBLEMS

Number of neutrons ϭ mass number Ϫ atomic number

Have students refer to Appendix
D for complete solutions to
Practice Problems.
14.

Number of neutrons ϭ 22 Ϫ 10 ϭ 12
Use the element name and mass number to write the isotope’s name. neon-22 Use the chemical symbol, mass number, and atomic number to write out the isotope in symbolic notation form.
22
10 Ne

Protons Neutrons and electrons

3. Evaluate the Answer
The relationships among number of electrons, protons, and neutrons have been applied correctly. The isotope’s name and symbol are in the correct format.

e!
Practic

PRACTICE PROBLEM
14. Determine the number of protons, electrons, and neutrons for isotopes b. through f. in the table above. Name each isotope, and write its symbol.

For more practice with problems using atomic number and mass number, go to
Supplemental Practice
Problems in Appendix A.

4.3 How Atoms Differ

101

Portfolio
Portfolio
Separating Isotopes
Linguistic Have students research how isotopes are separated for commercial, medical, and industrial purposes.
For instance, in nuclear power plants, only the uranium-235 is suitable for use as a fis-

sionable fuel and it must be separated from the more abundant uranium-238. Students may wish to diagram the sequence in addition to their writing. L2 P

Isotope

Symbol

b.

20

26

calcium-46

46
20Ca

c.

8

9

oxygen-17

17
8O

d.

26

31

iron-57

57
26Fe

e.

30

34

zinc-64

64
30Zn

f.

80

124

mercury-204

204
80Hg

Resource
Manager
ChemLab and MiniLab
Worksheets, p. 13 L2
Challenge Problems, p. 4 L3
Math Skills Transparency 4 and
Master L2 ELL
Teaching Transparency 13 and
Master L2 ELL

101

Mass of Individual Atoms

mini LAB
Purpose

Students will determine the atomic mass of a penny given a mixture of its pre- and post-1982 “isotopes.”
Process Skills

Classifying, measuring, using numbers Teaching Strategies

• Review the reason for the change

Table 4-2
Masses of Subatomic
Particles
Particle

Mass (amu)

Electron

0.000 549

Proton

1.007 276

Neutron

1.008 665

in composition of the penny in
1982 during the pre-lab discussion (cost savings).
• Students may have difficulty calculating the percent abundance and mass contribution of each isotope.
Expected Results

Pre-1982 pennies have a greater mass than post-1982 pennies.
Mass of ten pre-1982 pennies ϭ 31.10 g
Average mass of a pre-1982 penny ϭ 3.11g
Mass of ten post-1982 pennies ϭ 25.48 g
Average mass of a post-1982 penny ϭ 2.55 g
The atomic mass depends on the mixture analyzed.
Analysis

1. The relative number of preand post-1982 pennies in the bag determines the percentage abundance of each group.
The two abundances must sum to 100%.
2. The atomic mass of a penny depends upon the mixture of pennies each student receives.
Sample data is shown here. mass contribution (pre-1982) ϭ (55.0%)(3.11 g) ϭ1.71 g mass contribution (post-1982) ϭ (45.0%)(2.55 g) ϭ 1.15 g atomic mass ϭ(1.71 g ϩ 1.15 g) ϭ 2.86 g
3. Atomic mass is dependent upon the relative abundance of each isotope. A different mixture of pennies would have a different atomic mass.
4. Masses of individual pennies will vary due to wear. Using an average mass of each penny minimizes the errors related to the variation.

102

Recall from Table 4-1 that the masses of both protons and neutrons are approximately 1.67 ϫ 10Ϫ24 g. While this is a very small mass, the mass of
1
ᎏ an electron is even smaller—only about ᎏ
1840 that of a proton or neutron.
Because these extremely small masses expressed in scientific notation are difficult to work with, chemists have developed a method of measuring the mass of an atom relative to the mass of a specifically chosen atomic standard.
That standard is the carbon-12 atom. Scientists assigned the carbon-12 atom a mass of exactly 12 atomic mass units. Thus, one atomic mass unit (amu) is defined as ᎏ112ᎏ the mass of a carbon-12 atom. Although a mass of 1 amu is very nearly equal to the mass of a single proton or a single neutron, it is important to realize that the values are slightly different. As a result, the mass of silicon-30, for example, is 29.974 amu, and not 30 amu. Table 4-2 gives the masses of the subatomic particles in terms of amu.
Because an atom’s mass depends mainly on the number of protons and neutrons it contains, and because protons and neutrons have masses close to 1 amu, you might expect the atomic mass of an element to always be very near a whole number. This, however, is often not the case. The explanation involves how atomic mass is defined. The atomic mass of an element is the weighted

miniLAB
Modeling Isotopes
Formulating Models Because they have different compositions, pre- and post-1982 pennies can be used to model an element with two naturally occurring isotopes. From the penny "isotope" data, the mass of each penny isotope and the average mass of a penny can be determined.

Materials bag of pre- and post-1982 pennies, balance Procedure
1. Get a bag of pennies from your teacher, and sort the pennies by date into two groups: pre1982 pennies and post-1982 pennies. Count and record the total number of pennies and the number of pennies in each group.
2. Use the balance to determine the mass of ten pennies from each group. Record each mass to the nearest 0.01 g. Divide the total mass of each group by ten to get the average mass of a pre- and post-1982 penny "isotope."

Analysis
1. Using data from step 1, calculate the percentage abundance of each group. To do this, divide the number of pennies in each group by the total number of pennies.
2. Using the percentage abundance of each "isotope" and data from step 2, calculate the

102

Chapter 4 The Structure of the Atom

Assessment
Performance Have student groups design and carry out an experiment to determine the average mass of “isotopes” of an object such as beans or pasta. Use the
Performance Task Assessment List of
Designing an Experiment in PASC, p. 23.

atomic mass of a penny. To do this, use the following equation for each "isotope." mass contribution = (% abundance)(mass)
Sum the mass contributions to determine the atomic mass.
3. Would the atomic mass be different if you received another bag of pennies containing a different mixture of pre- and post-1982 pennies? Explain.
4. In step 2, instead of measuring and using the mass of a single penny of each group, the average mass of each type of penny was determined. Explain why.

Calculating the Weighted Average Atomic Mass of Chlorine
17eϪ

35
17

CI

17eϪ

Atomic mass: 34.969 amu
Percent abundance: 75.770%
17pϩ
Mass contribution:
(34.969 amu) (75.770%) ϭ 26.496 amu 20n0

17pϩ
18n0

37
17

Reinforcement

CI

Atomic mass: 36.966 amu
Percent abundance: 24.230%
Mass contribution:
(36.966 amu) (24.230%) ϭ 8.957amu

Enrichment

Weighted average atomic mass of chlorine ϭ (26.496 amu ϩ 8.957 amu) ϭ 35.453 amu

Visual-Spatial The mass

Figure 4-17

average mass of the isotopes of that element. For example, the atomic mass of chlorine is 35.453 amu. Chlorine exists naturally as a mixture of about 75% chlorine-35 and 25% chlorine-37. Because atomic mass is a weighted average, the chlorine-35 atoms, which exist in greater abundance than the chlorine-37 atoms, have a greater effect in determining the atomic mass. The atomic mass of chlorine is calculated by summing the products of each isotope’s percent abundance times its atomic mass. See Figure 4-17. For handson practice in calculating atomic mass, do the miniLAB on the previous page.
You can calculate the atomic mass of any element if you know its number of naturally occurring isotopes, their masses, and their percent abundances. The following Example Problem and Practice Problems will provide practice in calculating atomic mass.

To determine the weighted average atomic mass of chlorine, the mass contribution of each of the two isotopes is calculated, and then those two values are added together.

EXAMPLE PROBLEM 4–3
Isotope Abundance for
Element X

Calculating Atomic Mass
Given the data in the table at the right, calculate the atomic mass of unknown element X. Then, identify the unknown element, which is used medically to treat some mental disorders.
1. Analyze the Problem
You are given the data in the table. Calculate the atomic mass by multiplying the mass of each isotope by its percent abundance and summing the results. Use the periodic table to confirm the calculation and identify the element.
Known

Unknown

For isotope 6X: mass ϭ 6.015 amu abundance ϭ 7.50% ϭ 0.0750
For isotope 7X: mass ϭ 7.016 amu abundance ϭ 92.5% ϭ 0.925

atomic mass of X ϭ ? amu name of element X ϭ ?

Point out to students that the answer to a weighted average atomic mass problem will probably be a mass that is closest to the element with the highest percent abundance. Isotope

Mass
(amu)

Percent abundance 6X

6.015

7.5%

7X

7.016

92.5%

spectrometer is an instrument that can be used to separate the isotopes of a particular element.
A sample of the element is introduced into a vacuum and reduced to low pressure. Next, the vapor is passed through a beam of electrons, where ionization occurs. The ions produced are accelerated into a magnetic field. A charged particle traveling at high speeds follows a curved path whose radius depends on the speed of the particle and its mass-to-charge ratio (m/e). As a result of the differences of massto-charge, each isotope arrives at a different location on the collector plate, where a detector counts and displays a distribution of isotopes registered. This is called a mass spectrum. Have students research a mass spectrometer, sketching and labeling its main parts, and creating a flow chart of its function. L3

PROBLEMS
Have students refer to Appendix
D for complete solutions to
Practice Problems.
15. 10.81 amu
16. Helium-4 is more abundant

2. Solve for the Unknown
Calculate each isotope’s contribution to the atomic mass.
For 6X: Mass contribution ϭ (mass)(percent abundance) mass contribution ϭ (6.015 amu)(0.0750) ϭ 0.451 amu
Continued on next page

4.3 How Atoms Differ

103

in nature because the atomic mass of naturally occurring helium is closer to the mass of helium-4
(~4 amu) than to the mass of helium-3 (~3 amu).
17. 24.31 amu

Internet Address Book
Note Internet addresses that you find useful in the space below for quick reference.

National Science Content Standards

Pages 102–103
UCP.1, UCP.2; A.1; B.1

103

3 Assess

For 7X: Mass contribution ϭ (mass)(percent abundance) mass contribution ϭ (7.016 amu)(0.925) ϭ 6.490 amu

Check for Understanding

Sum the mass contributions to find the atomic mass.

Logical-Mathematical

Atomic mass of X ϭ (0.451 amu ϩ 6.490 amu) ϭ 6.941 amu

Obtain or create a copy of a mass spectrograph and have students analyze the abundance of each isotope and calculate an average atomic mass. The Internet or reference texts may be good sources for obtaining a mass spectrograph. L2

Reteach
Compare the calculation of an average atomic mass with the calculation of a student’s grade using weighted grades. For instance, suppose a student knew his or her grade in chemistry was weighted in the following way:
50% tests; 40% lab reports; 10% homework. If the student averaged
85 on tests, 95 on lab reports, and
70 on homework, what would be his final grade? 87.5

Use the periodic table to identify the element.
The element with a mass of 6.941 amu is lithium (Li).
3. Evaluate the Answer
The result of the calculation agrees with the atomic mass given in the periodic table. The masses of the isotopes have four significant figures, so the atomic mass is also expressed with four significant figures.

PRACTICE PROBLEMS e! Practic

For more practice with atomic mass problems, go to Supplemental
Practice Problems in
Appendix A.

16. Helium has two naturally occurring isotopes, helium-3 and helium-4.
The atomic mass of helium is 4.003 amu. Which isotope is more abundant in nature? Explain.
17. Calculate the atomic mass of magnesium. The three magnesium isotopes have atomic masses and relative abundances of 23.985 amu
(78.99%), 24.986 amu (10.00%), and 25.982 amu (11.01%).

Analyzing an element’s mass can give you insight into what the most abundant isotope for the element may be. For example, note that fluorine (F) has an atomic mass that is extremely close to a value of 19 amu. If fluorine had several fairly abundant isotopes, it would be unlikely that its atomic mass would be so close to a whole number. Thus, you might conclude that virtually all naturally occurring fluorine is probably in the form of fluorine19 (199F). You would be correct, as 100% of naturally occurring fluorine is in the form of fluorine-19. While this type of reasoning generally works well, it is not foolproof. Consider bromine (Br), with an atomic mass of 79.904 amu.
With a mass so close to 80 amu, it seems likely that the most common bromine isotope would be bromine-80 ( 80
35Br). This is not the case, however.
Bromine’s two isotopes, bromine-79 (78.918 amu, 50.69%) and bromine-81
(80.917 amu, 49.31%), have a weighted average atomic mass of approximately
80 amu, but there is no bromine-80 isotope.

Extension
Have students do a scavenger hunt through the Table of Isotopes in the
CRC Handbook of Chemistry and
Physics. Provide a list of clues
(give an element with the highest number of isotopes; three isotopes; a weighted average of 24.305) and a prize for the student who arrives at the correct element the fastest.

Section

L2

4.3

Assessment

18.

Which subatomic particle identifies an atom as that of a particular element? How is this particle related to the atom’s atomic number?

19.

What is an isotope? Give an example of an element with isotopes.

20.

Explain how the existence of isotopes is related to atomic masses not being whole numbers.

Assessment
Skill Pass around a sandwich bag that contains several varieties—or “isotopes”—of the same thing. Good choices are pastas, beans, and M&M candies.
Have students calculate the weighted average mass of the
“isotopes” present. L2

15. Boron has two naturally occurring isotopes: boron-10 (abundance ϭ
19.8%, mass ϭ 10.013 amu), boron-11 (abundance ϭ 80.2%, mass ϭ
11.009 amu). Calculate the atomic mass of boron.

104

Chapter 4 The Structure of the Atom

Section 4.3

Pages 104–105
UCP.1, UCP.2; B.1, B.6

104

Thinking Critically Nitrogen has two naturally occurring isotopes, N-14 and N-15. The atomic mass of nitrogen is 14.007 amu. Which isotope is more abundant in nature? Explain.

22.

Communicating List the steps in the process of calculating average atomic mass given data about the isotopes of an element.

chemistrymc.com/self_check_quiz

Assessment

18. The proton. The number of protons

National Science Content Standards

21.

equals the atomic number.
19. Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons.
Carbon has the three isotopes: C-12,
C-13, and C-14.
20. Atomic masses aren’t whole numbers because they represent weighted aver-

ages of the masses of the isotopes of an element. 21. N-14 must be more abundant because the atomic mass of nitrogen, which is a weighted average, is closer to the mass of N-14 than to the mass of N-15.
22. First, multiply the mass of each isotope by its percent abundance; then, sum the mass contributions of all the isotopes.

Section

4.4

Unstable Nuclei and Radioactive
Decay

Section 4.4

1 Focus
You now have a good understanding of the basic structure of matter and how matter interacts and changes through processes called chemical reactions.
With the information you have just learned about the atom’s nuclear nature, you are ready to learn about a very different type of reaction—the nuclear reaction. This section introduces you to some of the changes that can take place in a nucleus; you will revisit and further explore this topic in Chapter 25 when you study nuclear chemistry.

Objectives

Radioactivity

Vocabulary

• Characterize alpha, beta, and gamma radiation in terms of mass and charge.

nuclear reaction radioactivity radiation radioactive decay alpha radiation alpha particle nuclear equation beta radiation beta particle gamma ray

ELL

Section
Focus

16

Radiation
4, Sectio

Inc.

Chapter

What do es this sym bol mean
?
What sh ould you do when you see this sym bol? Matter and
Change

Section

c

4.4 Unstable Nuclei and Radioactive Decay

105

CHEMISTRY JOURNAL
Nuclear News
CD-ROM
Chemistry: Matter and Change
Demonstration: Evidence for
Alpha Particles

1
2

Chemistry:

Unstable systems, such as this pencil momentarily standing on its tip, gain stability by losing energy. In this case, the pencil loses gravitational potential energy as it topples over.
Unstable atoms also gain stability by losing energy—they lose energy by emitting radiation.

b

Transpare ncy Use with

Figure 4-18

a

Before presenting the lesson, display
Section Focus Transparency 16 on the overhead projector. Have students answer the accompanying questions using Section Focus
Transparency Master 16. L1

Copyright
© Glencoe/Mc
Graw-Hill,
a division of the McGr aw-Hill Comp anies, Recall from Chapter 3 that a chemical reaction involves the change of one or more substances into new substances. Although atoms may be rearranged, their identities do not change during the reaction. You may be wondering why atoms of one element do not change into atoms of another element during a chemical reaction. The reason has to do with the fact that chemical reactions involve only an atom’s electrons—the nucleus remains unchanged.
As you learned in the previous section, the number of protons in the nucleus determines the identity of an atom. Thus, because there are no changes in the nuclei during a chemical reaction, the identities of the atoms do not change. There are, however, reactions that do involve an atom of one element changing into an atom of another element. These reactions, which involve a change in an atom’s nucleus, are called nuclear reactions.
In the late 1890s, scientists noticed that some substances spontaneously emitted radiation in a process they called radioactivity. The rays and particles emitted by the radioactive material were called radiation. Scientists studying radioactivity soon made an important discovery—radioactive atoms undergo significant changes that can alter their identities. In other words, by emitting radiation, atoms of one element can change into atoms of another element. This discovery was a major breakthrough, as no chemical reaction had ever resulted in the formation of new kinds of atoms.
Radioactive atoms emit radiation because their nuclei are unstable.
Unstable systems, whether they’re atoms or the pencil standing on its sharpened tip shown in Figure 4-18a, gain stability by losing energy. As you can see in Figure 4-18b and Figure 4-18c, the pencil gains stability (and loses energy) by toppling over. When resting flat on the table top, the pencil has

Focus Transparency

• Explain the relationship between unstable nuclei and radioactive decay.

Have students scan through daily newspapers and magazines for articles dealing with nuclear phenomena. Have them keep a listing of the articles and classify them by categories, such as weaponry, power, food, medicine, and industrial. Based on their lists, have students estimate rough percentages for each category. L2

Focus Tran sparenc ies

2 Teach
Concept Development
Emphasize to students the three major kinds of changes: physical, chemical, and nuclear. Students should be able to describe the mechanism of each change and give clear examples of each.

Resource
Manager
Study Guide for Content
Mastery, p. 24 L2
Solving Problems: A Chemistry
Handbook, Section 4.4 L2
Section Focus Transparency 16 and Master L1 ELL
Teaching Transparency 14 and
Master L2 ELL

105

n 4.4

CAREERS USING CHEMISTRY

Radiation Protection
Technician

Career Path A career

Do you like the idea of protecting the health of others?
Would you enjoy the challenge of removing contaminated materials? If so, consider a career as a radiation protection technician. in handling hazardous materials would include high school courses in chemistry and math.
Technicians must complete high school and at least three months of training in handling nuclear waste, asbestos, and lead. They must also be licensed.
Career Issue Have students suggest reasons why radiation protection technicians are paid more than workers who remove lead or asbestos.

Radiation protection technicians use radiation measuring instruments to locate and assess the risk posed by contaminated materials. Then they decontaminate the area with high-pressure cleaning equipment and remove the radioactive materials. The work often requires the technicians to wear protective suits. The physically demanding work is carefully planned and carried out, with an emphasis on safety.

For More Information

For more information about careers in handling hazardous materials, students can contact:

Scientists began researching radioactivity in the late 1800s. By directing radiation from a radioactive source between two electrically charged plates, scientists were able to identify three different types of radiation. As you can see in Figure 4-19, some of the radiation was deflected toward the negatively charged plate, some was deflected toward the positively charged plate, and some was not deflected at all.
Alpha radiation Scientists named the radiation that was deflected toward the negatively charged plate alpha radiation. This radiation is made up of alpha particles. Each alpha particle contains two protons and two neutrons, and thus has a 2+ charge. As you know, opposite electrical charges attract.
So the 2+ charge explains why alpha particles are attracted to the negatively charged plate shown in Figure 4-19. An alpha particle is equivalent to a helium-4 nucleus and is represented by 42He or ␣. The alpha decay of radioactive radium-226 into radon-222 is shown below.
0

radium-226

Figure 4-19

To acquaint students with the concept of radiation, purchase an inexpensive radon kit from a local hardware store and test your classroom radon levels.
These kits are sent through the mail for laboratory analysis.
Radon, a byproduct of the decay of uranium-238 in soils and building materials, is considered hazardous due to the fact that it decays into radioactive isotopes of polonium, bismuth, and lead. These heavy metal ions are not easily eliminated from the body. It is estimated that 10 000 to 20 000 lung cancer deaths are caused by radon gas exposure in the
United States each year.

Types of Radiation

226 Ra
88

Laborers—AGC Education and Training Fund
37 Deerfield Road
P.O. Box 37
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Quick Demo

less gravitational potential energy than it did in its upright position, and thus, it is more stable. Unstable nuclei lose energy by emitting radiation in a spontaneous process (a process that does not require energy) called radioactive decay. Unstable radioactive atoms undergo radioactive decay until they form stable nonradioactive atoms, often of a different element. Several types of radiation are commonly emitted during radioactive decay.

Because alpha, beta, and gamma radiation possess different amounts of electrical charge, they are affected differently by an electric field. Gamma rays, which carry no charge, are not deflected by the electric field.

radon-222

4 He
2

alpha particle

Note also that a new element, radon (Rn), is created as a result of the alpha decay of the unstable radium-226 nucleus. The type of equation shown above is known as a nuclear equation because it shows the atomic number and mass number of the particles involved. It is important to note that both mass number and atomic number are conserved in nuclear equations. The accounting of atomic numbers and mass numbers below shows that they are conserved.
4
0 222
86 Rn ϩ 2 He
Atomic number: 88 0 86 ϩ 2
Mass number: 226 0 222 ϩ 4
226 Ra
88

Positive plate
Lead block

ϩ

222 Rn
86

Hole

Beta particles (1Ϫ charge)

␤

ϩ
Gamma rays
(no charge)

␥

Ϫ

Radioactive source Negative plate

Alpha particles (2ϩ charge)

␣

Zinc sulfide coated screen

106

Chapter 4 The Structure of the Atom

DIFFERENTIATED INSTRUCTION
Visually Impaired

National Science Content Standards

Pages 106–107
UCP.1, UCP.2; B.1, B.6; F.5

106

Kinesthetic Give students sets consisting of two differently shaped objects representing protons and neutrons. Use small containers, such as petri dishes, to represent the nucleus and contain the protons

and neutrons. Ask students to model various isotopes and then follow up by asking them to demonstrate an alpha or beta decay by removing the appropriate numbers of protons or neutrons. L1 ELL

Beta radiation Scientists named the radiation that was deflected toward the positively charged plate beta radiation. This radiation consists of fast moving electrons called beta particles. Each beta particle is an electron with a
1– charge. The negative charge of the beta particle explains why it is attracted to the positively charged plate shown in Figure 4-19. Beta particles are represented by the symbol Ϫ10 ␤. The beta decay of radioactive carbon-14 into nitrogen-14 is shown below.
0

14C
6

carbon-14

14N
7

+ nitrogen-14 0
Ϫ1␤

beta particle

Table 4-3
Characteristics of Alpha, Beta, and Gamma Radiation
Radiation Symbol Mass Charge type (amu)
Alpha

4 He
2

Beta

0
Ϫ1␤

Gamma

4

2ϩ

1
ᎏᎏ
1840

1–

0

0

0␥
0

The beta decay of unstable carbon-14 results in the creation of the new atom nitrogen (N).
Gamma radiation The third common type of radiation is called gamma radiation, or gamma rays. Gamma rays are high-energy radiation that possess no mass and are denoted by the symbol 00␥. Because they possess no electrical charge, gamma rays are not deflected by electric or magnetic fields.
Gamma rays usually accompany alpha and beta radiation, and they account for most of the energy lost during the radioactive decay process. For example, gamma rays accompany the alpha decay of uranium-238.
238U
92

0

uranium-238

4He ϩ ϩ
2 00␥
2
thorium-234 alpha particle gamma rays
234Th
90

Because gamma rays are massless, the emission of gamma rays by themselves cannot result in the formation of a new atom. Table 4-3 summarizes the basic characteristics of alpha, beta, and gamma radiation.
Nuclear stability You may be wondering why some atoms are stable while others are not. The primary factor in determining an atom’s stability is its ratio of neutrons to protons. The details of how the neutron-to-proton ratio determines stability will be covered in Chapter 25. For now, it is enough that you know that atoms containing either too many or too few neutrons are unstable. Unstable nuclei lose energy through radioactive decay in order to form a nucleus with a stable composition of neutrons and protons. Alpha and beta particles are emitted during radioactive decay, and these emissions affect the neutron-to-proton ratio of the newly created nucleus. Eventually, radioactive atoms undergo enough radioactive decay to form stable, nonradioactive atoms.
This explains why there are so few radioactive atoms found in nature—most of them have already decayed into stable atoms.

Section

4.4

Explain how unstable atoms gain stability. What determines whether or not an atom is stable?

24.

Create a table comparing the mass and charge of alpha, beta, and gamma radiation.

25.

In writing a balanced nuclear equation, what must be conserved?

26.

Thinking Critically Explain how a nuclear reaction differs from a chemical reaction.

chemistrymc.com/self_check_quiz

Section 4.4

27.

energy (emitting radiation). Atomic stability is determined primarily by the ratio of the number of neutrons to the number of protons.

Interpersonal Have student

groups of three use the Table of Isotopes in the CRC Handbook of Chemistry and Physics to find any isotope that undergoes radioactive decay. The first student in the group writes down the listed isotope, then passes the paper to a second student in the group. The second student writes down the ejected particle listed, then passes it to the third student in the group.
This student should complete the nuclear equation by writing the resulting nuclide. Have students continue with this process for five minutes. Each group should check with the teacher to see whether its equations are correct. The group with the most correct nuclear equations wins. L2 COOP LEARN

Have students complete the following nuclear reactions.
60
60
0
27Co → 28Ni ϩ ? Ϫ1␤
241
237
4
95Am → 93Np ϩ ? 2He

Extension
Have students research the use of radioactive isotopes in classical experiments in chemistry and physics. L2

Classifying Classify each of the following as a chemical reaction, a nuclear reaction, or neither.
Thorium emits a beta particle.
b. Two atoms share electrons to form a bond.
c. A sample of pure sulfur emits heat energy as it slowly cools.
d. A piece of iron rusts.
a.

4.4 Unstable Nuclei and Radioactive Decay

107

Assessment

23. Atoms can gain stability by losing

Check for Understanding

Reteach

Assessment

23.

3 Assess

24. Particle Symbol Mass (amu) Charge
␣

4
2He

␤

0
Ϫ1␤

␥

0
0␥

4

ϩ2

1/1840

Ϫ1

0

0

Assessment
Knowledge Have each student write two questions from each section of this chapter. Then use the questions for a game show similar to a quiz tournament. L2

26. Nuclear reactions involve changes in the nuclei of the atoms, usually resulting in changes of atoms of one element into atoms of another element. Chemical reactions involve changes of substances into new substances. 27. a. nuclear
c. neither
b. chemical
d. chemical

25. In a balanced nuclear equation, both mass number and atomic number must be conserved.

107

CHEMLAB

4

Preparation

CHEMLAB

4

Very Small Particles

Time Allotment

T

One laboratory period
Process Skills

Observing and inferring, formulating models, hypothesizing, comparing and contrasting
Safety Precautions

To avoid splashing your eyes or clothing with vanilla, do not deflate the balloon until the vanilla has dried inside.
Disposal

his laboratory investigation will help you conceptualize the size of an atom. You will experiment with a latex balloon containing a vanilla bean extract. Latex is a polymer, meaning that it is a large molecule (a group of atoms that act as a unit) that is made up of a repeating pattern of smaller molecules. The scent of the vanilla extract will allow you to trace the movement of its molecules through the walls of the solid latex balloon.

Problem

Objectives

Materials

How small are the atoms that make up the molecules of the balloon and the vanilla extract? How can you conclude the vanilla molecules are in motion?

• Observe the movement of vanilla molecules based on detecting their scent.
• Infer what the presence of the vanilla scent means in terms of the size and movement of its molecules.
• Formulate models that explain how small molecules in motion can pass through an apparent solid.
• Hypothesize about the size of atoms that make up matter. vanilla extract or flavoring 9-inch latex balloon (2) dropper The deflated balloons can be placed in a solid waste container.
Alternative Materials

After-shave lotion or perfume can be used to replace the vanilla.

Pre-Lab

Safety Precautions

2. A polymer is a large molecule made up of many repeating smaller molecules. Latex is an example. 3. The amount of vanilla placed into each balloon is constant.
The chemical composition of the balloons is constant.
4. The purpose of the vanilla flavoring is to trace the movement of molecules using the sense of smell to detect their presence. 5. The high-energy molecules leave first, lowering the average kinetic energy of the remaining liquid. It is cooler.
6. You are smelling a gas or vapor. • Always wear safety goggles and a lab apron.
• Be careful not to cut yourself when using a sharp object to deflate the balloon.

Pre-Lab
1.
2.
3.
4.
5.
6.
7.

Read the entire CHEMLAB.
Describe a polymer and give an example.
Identify constants in the experiment.
What is the purpose of the vanilla extract?
As a liquid evaporates, predict what you think will happen to the temperature of the remaining liquid.
When you smell an aroma, is your nose detecting a particle in the solid, liquid, or gas phase?
Prepare all written materials that you will take into the laboratory. Be sure to include safety precautions, procedure notes, and a data table in which to record your data and observations.

108

Data Table
Observations
Balloon 1 with vanilla

Initial

Final

Relative size
Relative
temperature

Balloon 2 without vanilla

Relative size
Relative
temperature

Chapter 4 The Structure of the Atom

Data Table
Observations
Balloon 1 with vanilla

Balloon 2 without vanilla

Initial

Final

Relative size Same size Same size Relative temp. Cool

Room temp. Relative size Same size Same size Relative temp. Room temp. Room temp. Procedure

Analyze and Conclude

• Encourage students to blow up the balloon so the walls are stretched tight. In this way, the pores of the walls become larger.
• Use large balloons (9-inch). The necks are larger, making it easier to add the vanilla; they are also easier to inflate and tie shut.

1. The volume of the balloons did not

Expected Results
108

See data table.

change.

2. Balloon 1 feels cool to the touch. The vanilla absorbs energy from the surroundings in changing to a vapor.
3. The vanilla’s odor is outside the balloon and fills the small, enclosed space. The vanilla vaporizes and its molecules move through the pores in the wall of the balloon. CHAPTER 4 CHEMLAB

Procedure
1.

Using the medicine dropper, add 25 to 30 drops of vanilla extract to the first balloon.

4. Vanilla will leak more rapidly

Cleanup and Disposal
After the vanilla has dried, deflate the balloon by puncturing it with a sharp object.
2. Dispose of the pieces of the balloon as directed by your teacher.
1.

Analyze and Conclude
1.

2.

3.

4.

2.

3.

4.

5.
6.

7.

Inflate the balloon so its walls are tightly stretched, but not stretched so tightly that the balloon is in danger of bursting. Try to keep the vanilla in one location as the balloon is inflated. Tie the balloon closed.
Feel the outside of the balloon where the vanilla is located and note the temperature of this area relative to the rest of the balloon. Record your observations in the data table.
Use only air to inflate a second balloon to approximately the same size as that of the first, and tie it closed. Feel the outside of the second balloon.
Make a relative temperature comparison to that of the first balloon. Record your initial observations.
Place the inflated balloons in a small, enclosed area such as a closet or student locker.
The next day, repeat the observations in steps 3 and 4 after the vanilla has dried inside the balloon.
Record these final observations.
To avoid splattering your clothes with dark brown vanilla, do not deflate the balloon until the vanilla has dried inside.

5.
6.

7.

Observing and Inferring How did the relative volumes of balloons 1 and 2 change after 24 hours? Explain.
Observing and Inferring By comparing the relative temperatures of balloons 1 and 2, what can you conclude about the temperature change as the vanilla evaporated? Explain.
Observing and Inferring Did the vanilla’s odor get outside the balloon and fill the enclosed space?
Explain.
Predicting Do you think vanilla will leak more rapidly from a fully inflated balloon or from a halfinflated balloon? Explain.
Hypothesizing Write a hypothesis that explains your observations.
Comparing and Contrasting Compare your hypothesis to Dalton’s atomic theory. In what ways is it similar? How is it different?
Error Analysis What factors might affect the results of different groups that performed the experiment? What types of errors might have occurred during the procedure?

Real-World Chemistry
Explain why helium-filled, Mylar-foil balloons can float freely for several weeks, but latex balloons for less than 24 hours.
2. How are high-pressure gases stored for laboratory and industrial use to prevent loss?
1.

from a fully inflated balloon.
The thinner wall will allow more vanilla to pass through its pores because the polymer chains are stretched farther apart.
5. A possible hypothesis is that the latex polymer chains move apart as the balloon inflates.
The stretching of the latex walls creates pores. The greater pressure inside the balloon causes the molecules to move from inside the balloon to the outside. The vanilla molecules are very small particles that fit through the pores and are detected by their aroma outside the balloon. The evaporation of the vanilla liquid requires energy. The wall of the balloon feels cool to the touch where the liquid is changing into a vapor. After 24 hours, the liquid evaporates and leaves a dried brown residue inside the balloon.
6. Dalton’s atomic theory states that all matter is composed of extremely small particles. So does the hypothesis in question 5. The laboratory investigation does not address the many other aspects of Dalton’s atomic theory.
7. Student answers will vary.
Temperature and inflation of the balloon (tightness) will affect results. The balloon may not have been tightly closed.

Real-World Chemistry
1. Metallic coated Mylar is a
CHEMLAB

Assessment
Performance Have a team of students repeat the laboratory investigation using small
Mylar balloons. Ask the students to make a hypothesis, in writing, about what they think will happen before repeating the investigation.
Use the Performance Task Assessment List for
Formulating a Hypothesis in PASC, p. 21. L2

109

Resource Manager

nonporous polymer film. Thus, helium does not escape Mylar.
2. High-pressure gases are stored in impermeable steel storage cylinders. ChemLab and MiniLab Worksheets, pp. 14–16 L2
National Science Content Standards

Pages 108–109
UCP.1, UCP.2; A.1; B.2, B.4

109

CHEMISTRY and

CHEMISTRY and

Purpose

Nanotechnology

Society

Society

Students will learn about nanotechnology and how it may affect the world in which they live.

Imagine a technology able to produce roads that repave themselves, greenhouses productive enough to end starvation, and computers the size of cells.
Sound like science fiction? It may not be.

Background

Starting at the bottom

Richard Feynmen, a scientist who won the 1965 Nobel prize in physics, first introduced the concept of nanotechnology in
1959. Since that time, many researchers have studied the possibilities. However, only recently, with the development of new and powerful scientific instruments and computer programs, is nanotechnology finally coming into reality.
Scientists now have the ability not only to visualize individual atoms, but to move them around as well.

This yet-to-be realized technology is generally known as nanotechnology. The prefix nano- means one billionth. A nanometer is roughly the size of several atoms put together. The goal of nanotechnology is to manipulate individual atoms in order to create a wide variety of products.
In order to manipulate the immense number of atoms required to make a product, scientists plan on constructing tiny robots called nanorobots. Nanorobots would have two objectives—to manipulate atoms and to copy themselves (self-replication).
Through self-replication, countless nanorobots could be created. This work force of nanorobots would then work together to quickly and efficiently assemble new products.

Teaching Strategies

The benefits of nanotechnology could potentially go beyond those of all other existing technologies. The quality and reliability of manufactured products could improve dramatically. For example, a brick could repair itself after cracks form, and a damaged road could repave itself. Furthermore, even as the quality and capability of products increase, their prices would decrease. With the use of nanorobot workforces and a readily available supply of atoms, the cost of atom-assembled products would be low.
Unlike many technological advances of the past, nanotechnology could also benefit the environment.
The need for traditional raw materials, such as trees and coal, would be greatly reduced. Also, because the atoms are arranged individually, the amount of waste could be carefully controlled and limited.
Thin materials called nanomembranes may even be able to filter existing pollutants out of air and water.
Advances in medicine could also be amazing.
Nanodevices smaller than human cells could be used to detect health problems, repair cells, and carry medicines to specific sites in the body.

Possibilities and effects

• Ask students to research the roles of enzymes, proteins, and ribosomes in the human body. Have them prepare posters, individually or in groups, in which they relate the roles of nanocomputers and nanorobots to naturally occurring processes in the human body.
• Ask students to choose an area that will be changed by nanotechnology, such as manufacturing, medicine, or environmental science. Have them consider what advances they would like to see in that area.
• Ask students to give an example of a revolutionary product that has been developed during their lifetimes. Have them explain how it has affected their lives.

Investigating the Issue
1. The Industrial Revolution began in the eighteenth century in England and later spread to other parts of the world. Students should describe several advances in technology that occurred during the period. They should also mention how these advances

110

110

When will it happen?
The promise of nanotechnology is still years away.
Some think nanodevices will be available in the next decade, whereas others expect the technology to take much longer to develop. So far, researchers have accomplished rearranging atoms into specific shapes such as letters and symbols, and have also succeeded in developing several very simple nanodevices. One such device, a nanoguitar, made of crystalline silicon, is about 1/20th the width of a single human hair. Another device under development will enable delivery of anti-cancer drugs to just the cancerous cells, while leaving normal cells unharmed. Investigating the Issue
1.

Communicating Ideas Read about the
Industrial Revolution and write a brief essay describing how technological advances affected society.

2.

Debating the Issue Nanotechnology supporters argue that dangers posed by the technology will be addressed as nanotechnology is developed. Should researchers be prevented from developing nanotechnology?

Visit the Chemistry Web site at chemistrymc.com to find links to more information about nanotechnology and the issues that surround it.

Chapter 4 The Structure of the Atom

lead to changes in socioeconomic structure and culture as society changed from agrarian to industrial. Finally, students should balance the positive results with a mention of the negative impact, such as increased pollution.
2. Answers will vary. Make sure that students back up their opinions with examples and facts. For example, one student might argue that nuclear technology offered many positive benefits

but it has been used for horrible destruction and still poses military dangers. For this reason, they might oppose the development of the technology. Another might argue that the automobile poses dangers to both drivers and pedestrians.
Yet laws and new technology, such as airbags, have been developed along the way to make driving as safe as possible.
They might argue that similar developments would make nanotechnology safe.

CHAPTER

4

STUDY GUIDE

CHAPTER
STUDY GUIDE

4

Using the Vocabulary

Summary
4.1 Early Theories of Matter
• The Greek philosopher Democritus was the first person to propose the existence of atoms.
• In 1808, Dalton proposed his atomic theory, which

was based on numerous scientific experiments.
• All matter is composed of atoms. An atom is the

smallest particle of an element that maintains the properties of that element. Atoms of one element are different from atoms of other elements.

• Atoms have equal numbers of protons and electrons,

and thus, no overall electrical charge.
• An atom’s mass number is equal to its total number

of protons and neutrons.
• Atoms of the same element with different numbers

of neutrons and different masses are called isotopes.
• The atomic mass of an element is a weighted aver-

age of the masses of all the naturally occurring isotopes of that element.

4.2 Subatomic Particles and the Nuclear Atom
• Atoms are composed of negatively charged electrons, neutral neutrons, and positively charged protons. Electrons have a 1Ϫ charge, protons have a 1ϩ charge, and neutrons have no charge. Both protons and neutrons have masses approximately 1840 times that of an electron.

4.4 Unstable Nuclei and Radioactive Decay
• Chemical reactions involve changes in the electrons surrounding an atom. Nuclear reactions involve changes in the nucleus of an atom.

• The nucleus of an atom contains all of its positive

• Alpha particles are equivalent to the nuclei of

charge and nearly all of its mass.
• The nucleus occupies an extremely small volume of

space at the center of an atom. Most of an atom consists of empty space surrounding the nucleus through which the electrons move.
4.3 How Atoms Differ
• The number of protons in an atom uniquely identifies an atom. This number of protons is the atomic number of the atom.

• The neutron-to-proton ratio of an atom’s nucleus

determines its stability. Unstable nuclei undergo radioactive decay, emitting radiation in the process. helium atoms, and are represented by 42 He or ␣.
Alpha particles have a charge of 2ϩ.
• Beta particles are high-speed electrons and are rep-

resented by Ϫ10 ␤. Beta particles have a 1Ϫ charge.

Review Strategies
• Have students summarize how the atomic model has changed from Dalton’s concept to present understanding. L2
• Have students list the parts of the atom and give their charge, location, and relative size. L2
• Have students determine the average atomic mass for an element given isotope masses and corresponding percent abundance. L2
• Problems from Appendix A or the
Supplemental Problems booklet can be used for review. L2

• Gamma rays are high-energy radiation and are rep-

resented by the symbol 00␥. Gamma rays have no electrical charge and no mass.

Key Equations and Relationships
• Determining the number of protons and electrons
Atomic ϭ number ϭ number number of protons of electrons
(p. 98)

To reinforce chapter vocabulary, have students write a sentence using each term. L2 ELL

• Determining the number of neutrons
Number
ϭ mass Ϫ atomic of neutrons number number

(p. 101)

For additional help with vocabulary, have students access the Vocabulary PuzzleMaker online. chemistrymc.com/ vocabulary_puzzlemaker Vocabulary
•
•
•
•
•
•
•
•

alpha particle (p. 106) alpha radiation (p. 106) atom (p. 90) atomic mass (p. 102) atomic mass unit (amu) (p. 102) atomic number (p. 98) beta particle (p. 107) beta radiation (p. 107)

•
•
•
•
•
•
•

cathode ray (p. 92)
Dalton’s atomic theory (p. 89) electron (p. 93) gamma ray (p. 107) isotope (p. 100) mass number (p. 100) neutron (p. 96)

chemistrymc.com/vocabulary_puzzlemaker

•
•
•
•
•
•
•

nuclear equation (p. 106) nuclear reaction (p. 105) nucleus (p. 95) proton (p. 96) radiation (p. 105) radioactive decay (p. 106) radioactivity (p. 105)

Study Guide

111

Use ExamView ® Pro Testmaker CD-ROM to:
■
■
■
■

Create multiple versions of tests.
Create modified tests with one mouse click for struggling students.
Edit existing questions and add your own questions.
Build tests based on national curriculum standards.

National Science Content Standards

Pages 110–111
UCP.1, UCP.2; B.1, B.2, B.4; E.1,
E.2; F.6; G.3

111

CHAPTER
CHAPTER

CHAPTER
ASSESSMENT

4

4
##

ASSESSMENT
ASSESSMENT

Concept Mapping

37. List the strengths and weaknesses of Rutherford’s

nuclear model of the atom. (4.2)

28. 1. matter; 2. atoms;
3. electrons; 4. protons or neutrons; 5. neutrons or protons; 6. empty space around nucleus; 7. nucleus

Mastering Concepts

32.

33.

34.
35.
36.
37.

38.
39.
40.
41.
42.

National Science Content Standards

Pages 112–113
UCP.1, UCP.2; B.1, B.2, B.4, B.6; G.3

112

What is the net charge of the nucleus? (4.2)
39. Explain what keeps the electrons confined in the space

surrounding the nucleus. (4.2)

Concept Mapping

40. Describe the flow of a cathode ray inside a cathode

28. Complete the concept map using the following terms:

electrons, matter, neutrons, nucleus, empty space around nucleus, protons, and atoms.

29. Democritus
30. John Dalton
31. The scientific instruments necessary to research matter at the atomic level had not been invented yet.
See Figure 4-4. Dalton was wrong about atoms being indivisible and all atoms of an element being identical.
Mass is conserved because atoms cannot be created, divided, or destroyed.
Chemical reactions involve only the separation, combination, and rearrangement of atoms. the electron changing the type of electrode or the type of gas did not affect the ray produced charge ϭ Ϫ1; mass ϭ 9.11 ϫ 10Ϫ28 g
Strengths: Rutherford’s model explained the results of the gold-foil experiment and why an atom is electrically neutral.
Weaknesses: The model could not account for the total mass of an atom or the arrangement of the electrons. protons and neutrons; net positive charge equal to the number of protons attraction to the positively charged nucleus from cathode to anode the plum pudding model
The ␣ particles were deflected by the positively charged gold nuclei.

38. What particles are found in the nucleus of an atom?

Go to the Chemistry Web site at chemistrymc.com for additional
Chapter 4 Assessment.

ray tube. (4.2)
41. Which outdated atomic model could be likened to

chocolate chip cookie dough? (4.2)
42. What caused the deflection of the alpha particles in

All

Rutherford’s gold foil experiment? (4.2)
43. Which subatomic particles account for most all of an

1.

atom’s mass? (4.2)

is composed of

44. How is an atom’s atomic number related to its number

2.

45. What is the charge of the nucleus of element 89? (4.2)

of protons? To its number of electrons? (4.2)
46. Explain why atoms are electrically neutral. (4.2)

which contain
3.

4.

47. Does the existence of isotopes contradict part of
5.

Dalton’s original atomic theory? Explain. (4.3)
48. How do isotopes of a given element differ? How are

in the
6.

in the
7.

they similar? (4.3)
49. How is the mass number related to the number of pro-

tons and neutrons an atom has? (4.3)
50. What do the superscript and subscript in the notation
40 K
19

Mastering Concepts
29. Who originally proposed the concept that matter was

composed of tiny indivisible particles? (4.1)
30. Whose work is credited with being the beginning of

modern atomic theory? (4.1)
31. Explain why Democritus was unable to experimentally

verify his ideas. (4.1)
32. State the main points of Dalton’s atomic theory using

your own words. Which parts of Dalton’s theory were later found to be in error? Explain why. (4.1)
33. Explain how Dalton’s atomic theory offered a con-

vincing explanation of the observation that mass is conserved in chemical reactions. (4.1)
34. Which subatomic particle was discovered by

researchers working with cathode ray tubes? (4.2)
35. What experimental results led to the conclusion that

electrons were part of all forms of matter? (4.2)
36. What is the charge and mass of a single electron? (4.2)

112

Chapter 4 The Structure of the Atom

43. protons and neutrons
44. They are all equal.

represent? (4.3)

51. Explain how to determine the number of neutrons an

atom contains if you know its mass number and its atomic number. (4.3)
52. Define the atomic mass unit. What were the benefits

of developing the atomic mass unit as a standard unit of mass? (4.3)
53. What type of reaction involves changes in the nucleus

of an atom? (4.4)
54. Explain how energy loss and nuclear stability are

related to radioactive decay. (4.4)
55. Explain what must occur before a radioactive atom

ceases to undergo further radioactive decay. (4.4)
56. Write the symbols used to denote alpha, beta, and

gamma radiation and give their mass and charge. (4.4)
57. What change in mass number occurs when a radioac-

tive atom emits an alpha particle? A beta particle? A gamma particle? (4.4)
58. What is the primary factor determining whether or not

an atom is stable or unstable? (4.4)

chemistrymc.com/chapter_test

Resource Manager
Chapter Assessment, pp. 19–24 L2
Supplemental Problems, Ch. 4
ExamView® Pro CD-ROM
MindJogger Videoquizzes DVD/VHS
Solutions Manual, Ch. 4
Chemistry Interactive CD-ROM,
Ch. 4 quiz

CHAPTER 4 ASSESSMENT
CHAPTER 4 ASSESSMENT

Mastering Problems

Mixed Review

Atomic Number and Mass Number (4.3)

Sharpen your problem-solving skills by answering the following. 59. How many protons and electrons are contained in an

45. 89ϩ
46. The number of positively

atom of element 44?
60. For each of the following chemical symbols, deter-

mine the element name and the number of protons and electrons an atom contains.
a. V
b. Mn

c. Ir
d. S

69. Describe a cathode ray tube and how it operates.

47.

70. Explain how J. J. Thomson’s determination of the

charge-to-mass ratio of the electron led to the conclusion that atoms were composed of subatomic particles.
71. How did the actual results of Rutherford’s gold foil

61. A carbon atom has a mass number of 12 and an atomic

48.

experiment differ from the results he expected?
72. Complete the table below.

number of 6. How many neutrons does it have?
62. An isotope of mercury has 80 protons and 120 neu-

trons. What is the mass number of this isotope?

Composition of Various Isotopes

63. An isotope of xenon has an atomic number of 54 and

Isotope Atomic
Mass Number of
Number
Number of number number protons of neutrons electrons

contains 77 neutrons. What is the xenon isotope’s mass number?

32

50.

16

64. How many electrons, protons, and neutrons are con-

24

tained in each of the following atoms?
a.
b.

132 Cs
55
59 Co
27

c.
d.

20

Zn-64

163 Tm
69
70 Zn
30

9
11

65. How many electrons, protons, and neutrons are con-

10

51.

23

tained in each of the following atoms?
a. gallium-64
b. fluorine-23

52.

c. titanium-48
d. helium-8

73. Approximately how many times greater is the diame-

Atomic Mass (4.3)
66. Chlorine, which has an atomic mass of 35.453 amu,

ter of an atom than the diameter of its nucleus?
Knowing that most of an atom’s mass is contained in the nucleus, what can you conclude about the density of the nucleus?

has two naturally occurring isotopes, Cl-35 and Cl-37.
Which isotope occurs in greater abundance? Explain.

74. Is the charge of a nucleus positive, negative, or zero?

107 Ag
47

75. Why are electrons in a cathode ray tube deflected by

67. Silver has two isotopes,

has a mass of 106.905 amu (52.00%), and 109
47 Ag has a mass of 108.905 amu
(48.00%). What is the atomic mass of silver?

68. Data for chromium’s four naturally occurring isotopes

is provided in the table below. Calculate chromium’s atomic mass.

The charge of an atom? magnetic and electric fields?

Chromium Isotope Data
Percent abundance

Mass (amu)

Cr-50

4.35%

49.946

Cr-52

83.79%

51.941

Cr-53

9.50%

52.941

Cr-54

2.36%

53.939

53.
54.

76. What was Henry Moseley’s contribution to our under-

standing of the atom?
77. What is the mass number of potassium-39? What is

the isotope’s charge?
78. Boron-10 and boron-11 are the naturally occurring iso-

Isotope

49.

topes of elemental boron. If boron has an atomic mass of 10.81 amu, which isotope occurs in greater abundance?
79. Calculate the atomic mass of titanium. The five tita-

nium isotopes have atomic masses and relative abundances of 45.953 amu (8.00%), 46.952 amu (7.30%),
47.948 amu (73.80%), 48.948 amu (5.50%), and
49.945 amu (5.40%).

55.

charged protons equals the number of negatively charged electrons.
Yes; not all atoms of an element are identical in mass. differ: number of neutrons, masses; similar: chemical properties, number of protons and electrons mass number ϭ number of pϩϩ number of n0
The superscript represents the mass number (40) and the subscript represents the atomic number (19). number of n0 ϭ mass number – atomic number amu ϭ 1/12 of the mass of a
C-12 atom; Scientists defined the atomic mass unit as a relative standard that was closer in size to atomic and subatomic masses. nuclear reaction
Radioactivity results when unstable nuclei emit energy in order to gain stability. A stable, nonradioactive atom must be formed.

56.
Particle Symbol Mass (amu) Charge

␣

4

ϩ2

␤

0
Ϫ1␤

1/1840

Ϫ1

␥

0
0␥

0

4
2He

0

57. ␣, mass number decreases

Assessment

60.

Element

Protons

a.

vanadium

23

23

b.

manganese

25

25

c.

iridium

77

77

d.

sulfur

16

16

61. 6 neutrons
Level 2
62. mass number is 200
63. mass number is 131

Electrons

113

by 4; ␤, no change in mass number; ␥, no change in mass number
58. the neutron-to-proton ratio

Mastering Problems

64.

Symbol Electrons Protons Neutrons
a.
b.
c.
d.

65.

132
55Cs
59
27Co
163
69Tm
70
30Zn

55
27
69
30

55
27
69
30

77
32
94
40

Symbol Electrons Protons Neutrons
a.
b.
c.
d.

Ga-64
F-23
Ti-48
He-8

31
9
22
2

31
9
22
2

33
14
26
6

Complete solutions to Chapter
Assessment problems can be found in the Solutions Manual.
Atomic Number and Mass
Number (4.3)
Level 1
59. 44 protons, 44 electrons
113

CHAPTER
CHAPTER 4 ASSESSMENT

Atomic Mass (4.3)
Level 1
66. Cl-35 must be more abun-

ASSESSMENT

4

80. Identify the two types of radiation shown in the figure

below. Explain your reasoning.

dant because the atomic mass of chlorine is much closer to the mass Cl-35 than the mass of Cl-37.

Level 2
67. 107.86 amu
68. 51.99 amu

ϩ

89. Individual atoms can be seen using a sophisticated

device known as a scanning tunneling microscope.
Write a short report on how the scanning tunneling microscope works and create a gallery of scanning tunneling microscope images from sources such as books, magazines, and the Internet.
Ϫ

81. Describe how each type of radiation affects an atom’s

atomic number and mass number.

Mixed Review

82. Silicon is very important to the semiconductor manu-

69. A cathode ray tube has a metal electrode at each end and is filled with a gas at low pressure. One electrode is connected to the negative terminal of a battery
(cathode), and the other is connected to the positive terminal (anode). When current flows, negatively charged particles called electrons are emitted from the cathode and travel through the tube to the anode.
70. Thomson determined that the electron’s mass was much less than the mass of a hydrogen atom, the lightest atom. This showed that there were smaller, subatomic particles. Atoms are divisible.
71. Rutherford expected the ␣ particles to be slightly deflected when they passed through a thin gold foil. Instead, he found that some were deflected at very large angles.

facturing industry. The three naturally occurring isotopes of silicon are silicon-28, silicon-29, and silicon-30. Write the symbol for each.

Thinking Critically
83. Applying Concepts Which is greater, the number of

compounds or the number of elements? The number of atoms or the number of isotopes? Explain.
84. Analyzing Information An element has three natu-

rally occurring isotopes. What other information must you know in order to calculate the element’s atomic mass? 85. Applying Concepts If atoms are primarily com-

posed of empty space, why can’t you pass your hand through a solid object?
86. Formulating Models Sketch a modern atomic

model of a typical atom and identify where each type of subatomic particle would be located.
87. Applying Concepts Copper has two naturally occur-

ring isotopes and an atomic mass of 63.546 amu. Cu63 has a mass of 62.940 amu and an abundance of
69.17%. What is the identity and percent abundance of copper’s other isotope?

Writing in Chemistry
88. The Standard Model of particle physics describes all

of the known building blocks of matter. Research the particles included in the Standard Model. Write a short report describing the known particles and those thought to exist but not detected experimentally.

Cumulative Review
Refresh your understanding of previous chapters by answering the following.
90. How is a qualitative observation different from a

quantitative observation? Give an example of each.
(Chapter 1)
91. A 1.0-cm3 block of gold can be flattened to a thin

sheet that averages 3.0 ϫ 10-8 cm thick. What is the area (in cm2) of the flattened gold sheet? A letter size piece of paper has an area of 603 cm2. How many sheets of paper would the gold cover? (Chapter 2)
92. Classify the following mixtures as heterogeneous or

homogeneous. (Chapter 3)
a.
b.
c.
d.

salt water vegetable soup
14-K gold concrete 93. Are the following changes physical or chemical?

(Chapter 3)
a.
b.
c.
d.
e.

water boils a match burns sugar dissolves in water sodium reacts with water ice cream melts

72.
Isotope Atomic number Mass pϩ number

n0

eϪ

114

S-32

16

32

16

16

16

Ca-44

20

44

20

24

20

Zn-64

30

64

30

34

30

9

19

9

10

9

11

23

11

12

11

F-19
Na-23

73. An atom’s diameter is about
10 000 times the diameter of its nucleus. Consequently, the density of the nucleus must be enormous.

114

Chapter 4 The Structure of the Atom

74. The nucleus is positively charged, whereas the atom is neutral.
75. Because they are charged particles, and charged particles are affected by the electrostatic forces of attraction and repulsion from electric and magnetic fields.
76. Moseley discovered that each element contains a unique positive charge (or number of protons) in its nucleus.

Thus, the number of protons in an atom’s nucleus uniquely identifies it as an atom of a particular element.
77. mass number ϭ 39; charge ϭ 0
78. B-11 must occur in greater abundance because the atomic weight of bromine is much closer to the mass of
B-11 than to the mass of B-10.
79. 47.89 amu

STANDARDIZED TEST PRACTICE
CHAPTER 4

CHAPTER 4 ASSESSMENT

Use these questions and the test-taking tip to prepare for your standardized test.
1. An atom of plutonium
a. can be divided into smaller particles that retain all

the properties of plutonium.
b. cannot be divided into smaller particles that retain

all the properties of plutonium.
c. does not possess all the properties of a larger quantity of plutonium.
d. cannot be seen using current technology.
237 Np,
93
decays by emitting one alpha particle, one beta particle, and one gamma ray. What is the new atom formed from this decay?

2. Neptunium’s only naturally occurring isotope,

a.
b.

233 U
92
241 Np
93

c.
d.

233 Th
90
241 U
92

3. An atom has no net electrical charge because
a. its subatomic particles carry no electrical charges.
b. the positively charged protons cancel out the nega-

tively charged neutrons.
c. the positively charged neutrons cancel out the negatively charged electrons.
d. the positively charged protons cancel out the negatively charged electrons.
4.

126 Te
52

a.
b.
c.
d.

has

126 neutrons, 52 protons, and 52 electrons.
74 neutrons, 52 protons, and 52 electrons.
52 neutrons, 74 protons, and 74 electrons.
52 neutrons, 126 protons, and 126 electrons.

5. Assume the following three isotopes of element Q

exist: 248Q, 252Q, and 259Q. If the atomic mass of Q is
258.63, which of its isotopes is the most abundant?
a.
b.

248Q

259Q

c.
d. they are all equally

252Q

abundant

80. The deflected beam is ␣

6. Based on the table, what is the mass of an atom of

neon found in nature?
a.
b.
c.
d.

19.992 amu
20.179 amu
20.994 amu
21.991 amu

7. In which of the neon isotopes is the number of neu-

trons the same as the number of protons?
a. 20Ne
b. 21Ne
c. 22Ne
d. none of the above
8. The atomic mass of Ne is equal to _____ .

19.922 amu ϩ 20.994 amu ϩ 21.991 amu
3
1
b. ᎏᎏ[(19.992 amu)(90.48%) ϩ (20.994 amu)(0.27%) ϩ
3
(21.991 amu)(9.25%)]
a. ᎏᎏᎏᎏᎏ

c. (19.992 amu)(90.48%) ϩ (20.994 amu)(0.27%) ϩ

(21.991 amu)(9.25%)
d. 19.992 amu ϩ 20.994 amu ϩ 21.991 amu

radiation because it is deflected toward the negatively charged plate. The undeflected beam must be neutral ␥ gamma radiation.
81. ␣, atomic number decreases by 2, mass number decreases by 4; ␤, atomic number increases by 1, mass number unchanged; ␥, atomic number and mass number are unchanged
29
30
82. 28
14Si, 14Si, 14Si

Thinking Critically
83. The number of compounds

9. Element X has an unstable nucleus due to an over-

abundance of neutrons. All of the following are likely to occur EXCEPT
a. element X will undergo radioactive decay.
b. element X will eventually become a stable, nonra-

dioactive element.
c. element X will gain more protons to balance the

neutrons it possesses.
d. element X will spontaneously lose energy.
10. The volume of an atom is made up mostly of
a.
b.
c.
d.

protons. neutrons. electrons. empty space.

84.

Interpreting Tables Use the table to answer questions 6–8.

85.

Characteristics of Naturally Occurring Neon
Isotopes
Isotope

Atomic number Mass
(amu)

Percent abundance 20Ne

10

19.992

90.48

21Ne

10

20.994

0.27

22Ne

10

21.991

9.25

Skip Around If You Can

The questions on some tests start easy and get progressively harder, while other tests mix easy and hard questions. You may want to skip over difficult questions and come back to them later, after you’ve answered all the easier questions. This will guarantee more points toward your final score. In fact, other questions may help you answer the ones you skipped. Just be sure you fill in the correct ovals on your answer sheet.

chemistrymc.com/standardized_test

Cumulative Review
90. A qualitative observation does not involve measurement (water is hot). A quantitative observation involves measurement (the water is 42oC).
91. area ϭ 3.3 ϫ 107 cm2; sheets ϭ 55 000
92. a. homogeneous
b. heterogeneous
c. homogeneous
d. heterogeneous

Standardized Test Practice

93. a.
b.
c.
d.
e.

b a d b 87.
115

physical change chemical change physical change chemical change physical change

Standardized Test Practice
1.
2.
3.
4.

86.

is greater than the number of elements because compounds are combinations of elements and the elements can be combined in many ways. The number of isotopes is greater than the number of atoms because each element has only one type of atom but may have more than one isotope. You also need to know the mass and percent abundance of each isotope.
Atoms in a solid object are bonded together by electrical forces—bonds that are not easily broken.
These tightly bonded atoms form solid objects.
Sketches should look similar to Figure 4-13.
The other isotope is Cu-65.
Its percent abundance is
30.83%.

5. c
6. a
7. a

8. c
9. c
10. d

National Science Content Standards

Pages 114–115
UCP.1, UCP.2; B.1, B.6

115