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Atomic Structure
Goals
Describe Daiton'satomic theory and itssignificance in thestudy of matter.
Infer a conceptualmodel of the structureof an atom, includtngthe properties of themajor subatomicparticles.
Demonstrate therelationship betweenthe atomic mass of anelement and the iso-topes of that element.
Shown here is the interiorof a scanning tunnelingmicroscope (STM), themost powerful micro-scope In the world. Witha magnification of 100million, an STM can probematter at the atomic level,making atoms visible tothe human eyel
Concept Overview
AtomicStructure
electrons protons neutrons
The Concept Overview organizes atomicthe major concepts of this chapter, numberThis diagram shows one way tolink these concepts related to atomicstructure.
mass number
ave you ever been asked to believe in something you couldn'tsee? You cannot see the basic components of matter with your
eyes. Yet invisible atoms are the fundamental units of which all matteris composed. Even more remarkable, the atom itself can be fracturedinto many pieces.
In this chapter, you will enter the world of atomic structure. Youwill meet some interesting people. Democritus and John Dalton wereteachers separated in history by more than 2000 years. Not long afterDalton's life ended, Sir J. J. Thomson and Ernest Rutherford were born.Thomson and Rutherford were physicists who lived into this century.The story of these thinkers and experimenters is filled with ideas anddiscoveries that include the concept of an atom and the detection ofsubatomic particles.
4.1 AtomsDemocritus of Abdera, a famous teacher who lived in fourth centuryB.C. Greece, first suggested the idea of atoms. Democritus was partof the "Atomists" school of thought. The Atomists thought thatmatter was composed of tiny indivisible particles called atoms.Democritus said that atoms were invisible, indestructible, funda-mental units of matter. Democritus's ideas agreed with later
nucleus
isotopes
Objective
Summarize Dalton's atomictheory.
4.1 Atoms 83
Biographical
Note
John Dalton(1766-1844)
Dalton was only 12 years oldwhen he took his first job as aschool teacher, Throughouthis life, he earned his livingas a teacher. However,Dalton's real love and his realgenius were for science.During his lifetime, he studiedmany different topics, includingthe aurora borealis, the tradewinds, and color blindness,Dalton's most notable contri-bution to science was hisreintroduction of the idea ofatoms to explain chemicalbehavior. In spite of its flaws,Dalton's atomic theoryinspired a generation ofchemists to study atomictheory.
Ftguro 4.1 According to Dal-ton's atomic theory, an element isa large collection of atoms and acompound is a large collection ofmolecules.
Chapter 4 Atomic Structure
scientific theory. However, his ideas were not useful in explaining
chemical behavior because they lacked experimental support.
Democritus thought and talked about atoms, but scientific experi-
ments were unknown in his world. Therefore Democritus did no
experiments to test his theories. The connection between observ-
able chemical changes and events at the level of individual atoms
was not to be established for 2200 years.
An English school teacher, John Dalton (1766—1844), studied
chemistry very differently from Democritus, who only philoso-
phized about atoms. Unlike Democritus, Dalton performed
experiments to arrive at his atomic theory. Dalton wanted to learn
in what ratios different elements combine in chemical reactions.
Based on the results of his experiments, Dalton formulated
hypotheses and theories to explain his observations. Eventually he
devised Dalton's atomic theory, which includes the following ideas,
illustrated in Figure 4.1.
1. All elements are composed of submicroscopic indivisible parti-
cles called atoms.
2. Atoms of the same element are identical. The atoms of any one
element are different from those of any other element.
3. Atoms of different elements can physically mix together or canchemically combine with one another in simple whole-numberratios to form compounds.
4. Chemical reactions occur when atoms are separated, joined, orrearranged. However, atoms of one element are never changed intoatoms of another element as a result of a chemical reaction.
o
oO
Atoms ofelement A
o
OO
o
o(b) (c)
Atoms of Mixture ofelement B atoms of
elementsA and B
(d)Compound ofmolecules madeby uniting atoms
ot elementsA and B
An ordinary coin the size of a penny but composed of urecopper (Cu) illustrates Dalton's concept of the atom. Imayih t
you could grind the copper into fine dust. Each speck in your sn{all
pile of shiny red dust would still have the of copper.
Suppose you continued to divide the specks of coppe into smaller
parts. Eventually, you would come upon a particle fzcppper that
could no longer be divided and still have the prop r ie of copper.
ment that ms the ro ertiesCopper atoms are very small. Your hypo et cal ure copper
coin the size of a penny would contain abou 2.4 x atoms. By
comparison, the earth's population is ago 4 1 A pe ple. Ther
are about 6 x 10 12 as many atoms in yo lit ec there e
people on earth. Even a speck of coppéfd st nta•
ably large number of atoms.
Does seeing individual atom s e i p s i e? Despite their
small size, individual atoms arg •s' I wi th proper instrument.
A scanning tunneling microscop visu zes individual atoms, as
you can see in Figure 4.2. Individ I ato s can even be arranged in
patterns. The ability to move indi dual toms holds promise for the
future creation of atomic-size e ctr devices such as circuits
and computer chips. This atomic scal fechnology could be applied
to computers, communications, nd ace exploration.
Figure 4.2 The surface Ofindividual gold atoms appears in
this photograph taken with a
scanning tunneling microscope.
Roots of words
atg;nos: (Greek) indivisibletom the smallest particle ofn element that retains the
properties of that element
Atoms are far too small to seewith the unaided eye.
4.1 Atoms
have
42 tlcctrcns, Protons, and neutrons
ee oi the
n hgute
iev
of
etee current
the '[be elee
were connected to a high-voltage source of electricity. One elec-trode, the anode, became positively charged. The other electrode,the cathode, became negatively charged. A glowing beam formedbetween the electrodes. The glowing beam, which travels from thecathode to the anode, is a cathode ray.
Thomson found that cathode rays were attracted to metal platesthat carry a positive electrical charge. The rays were repelled byplates that carry a negative electrical charge. Figure 4.4 shows adeflection of the cathode rays. In electricity, opposite charges attractand like charges repel. Therefore Thomson proposed that a cathoderay is a stream of very small negatively charged particles, all alike,moving at high speed. He called these particles electrons. Moreover,Thomson showed that cathode rays are always composed of elec-trons, regardless of the kind of gas in the cathode ray tube or thekind of metal in the electrodes. Thomson concluded that electronsmust be a part of the atoms of all elements. By 1900, Thomson andothers had figured out that the electron carries roughly one unit ofnegative charge and that its mass is about 1/2000 the mass of ahydrogen atom.
Clever experiments enabled the American scientist Robert A.Millikan (1868—1953) to improve earlier estimates of the charge onan electron. Because he had accurate values of both the charge andthe ratio of the charge to the mass of an electron, Millikan could cal-culate an accurate value for the mass of the electron. Millikan'svalues of charge and mass, reported in 1916, are very similar tothose accepted today. An electron carries exactly one unit of nega-tive charge and its mass is 1/1840 the mass of a hydrogen atom.
The most common form of the hydrogen atom is the lightestatom that exists. If an electron is only 1/1840 the mass of a commonhydrogen atom, what is left over when one of these atoms losesan electron? You can think through this problem with four simpleideas about matter and electric charges. First, atoms have no elec-tric charge; they are electrically neutral. The evidence for electrical
High voltage
Negativelycharged plate
Cathode Screen with Direction of Positively Anodehole cathode ray charged plate
To vacuumpump
New Discoveries ofSubatomic ParticlesIn the decades following thediscovery of the neutron,physicists discovered moresubatomic particles.Physicists found that whenatomic nuclei were struckby high-energy particles orradiation, a target nucleuswould shatter, creating unsta-ble particles. This happenedinfrequently in nature, sophysicists designed andbuilt cyclotrons, synchro-trons, and linear acceleratorsto create collisions and studythem in detail. As a result ofthese tools, hundreds ofsubatomic particles havebeen discovered.
Figure 4.4 Cathode rays areattracted by a positively chargedplate. This attraction shows thenegatively charged character ofthe particles.
4.2 Electrons, Protons, and Neutrons 87
Figure 4.5 If the gas in thecathode ray tube was hydrogen,
High voltage
the canal rays would be made upof protons; after hydrogen gasatoms lose electrons at thecathode, only protons remain toform the canal ray.
Canal rays(positive particles)
Direction of Cathode canal rays
Tabte 4.1*'Properties of Subatomic Particles
Cathode ray(electrons) +
Screen with Direction of 1 Anodehole cathode ray
To vacuum pump
Approximate
Particle
Electron
Proton
Neutron
1 amu =
Symbol
10 -24 g.
Relativeelectrical charge
1+
relative mass(amu)*
1/18401
1
Actual mass(g)
9.11 x 10-28
1.67 x 10-24
1.67 x 10-24
Roots of words
protos: (Greek) firstproton a positively chargedsubatomic particle found inthe nucleus of an atom
The proton was the firstnuclear particle to bediscovered.
elektron: (Greek) shiningbeam
electron a negativelycharged atomic particle
The path taken by a swarmof electrons shows up as ashining beam in a cathoderay tube.
88 Chapter 4 Atomic Structure
neutrality is that you do not receive an electric shock every time youtouch an object. Second, electric charges are properties of particlesof matter. That is, electric charges are carried by particles of matter.Third, electric charges exist in a single unit or in multiples of a singleunit. There are no fractions of charges. Fourth, electric chargescancel when equal numbers of negatively charged and positivelycharged particles combine to form an electrically neutral particle.Thus, because an electron carries one unit of negative charge, thereshould be a particle with one unit of positive charge left over when acommon hydrogen atom loses an electron. This positively chargedsubatomic particle is called a proton. In 1886 E. Goldstein, using acathode ray tube in which the cathode had holes, observed raystraveling in the opposite direction to the cathode ray. These rays,shown in Figure 4.5, contain positively charged particles and arecalled canal rays.
In 1932 the English physicist Sir James Chadwick (1891—1974)confirmed the existence of yet another subatomic particle: the neu-tron. Neutrons are subatomic particles with eo charge, but theirmass nearly equals that of the proton. Thus the fundamentalbuilding blocks of atoms are the electron, the proton, and theneutron. Table 4.1 summarizes the properties of these three sub-atomic particles.
Concept Practice
3. Since all atoms have negatively charged electrons, shouldn'tevery sample of matter have a negative charge? Explain.
4. What experimental evidence did Thomson have for thefollowing ideas?a. Electrons have a negative charge.b. Atoms of all elements contain electrons.
Charge
To describe the properties ofelectrical charge.
four 25-cm pieces ofclear plastic tape
two round balloons
two 60-cm pieces of string
one piece of wool or
acrylic material
l. Put two 25-cm pieces of thetape on opposite sides of yoursmooth desk top, leaving 2 to 3cm sticking over the edge. Graspthe ends of the tape. Pull both ofthe pieces of tape from the deskand slowly bring them towardone another. What do you
2. You and your lab partnershould each take a third 25-cmpiece of tape and pull it betweentwo of your fingers (as if you weretrying to clean it). Slowly bringthis piece of tape together withthe fourth piece. What do youobserve?
3. What would you predict mighthappen if you brought a piece oftape pulled up from the desk topclose to a piece of tape pulledthrough your fingers? Try it! Whathappened?
4. Inflate both balloons and tie astring to each. Rub both balloonsagainst your hair. Hold the bal-loons by the strings and slowlymove them toward each other.What happens?
5. Rub the balloons against thepiece of material, and slowlymove them toward each other.Rub the balloons against yourown clothes and move themtogether. What happens?
Analysis and ConclusionsI. How do objects with likecharges react to one another?What about objects with oppo-site charges?
2. What subatomic particle istransferred when objects becomeelectrically charged?
3. How does an object becomepositively charged?
4. Can you predict whether apiece of tape pulled from a desktop would be attracted orrepelled by a balloon rubbed onwool? Explain.
5. If you knew how the tapepulled from the desk top reactedto the charged balloon, couldyou predict how the tape pulledthrough your fingers would reactwith the balloon? Why?
observe?
4.2 Electrons, Protons, and Neutrons 89
Objective
Explain the structure of anatom including the location ofthe proton, electron, andneutron With respect to thenucleus.
(a) Rutherford andMarsden aimed a beam of alphaparticles at a piece cf gold foilsurrounded by a fluorescentscreen. (b) Only particles thatpass near or approach thenucleus directly are affected.
90 Chapter 4 Atomic Structure
4.3 The structure of theNuclear AtomEven before neutrons were discovered, scientists were wondering
how electrons and protons were put together in an atom. The pre-
vailing theory was that the protons and electrons were evenly
distributed throughout the volume of an atom. In 1911, Ernest
Rutherford (1871—1937) and his co-workers at the University of
Manchester, England, decided to test this theory of atomic struc-
ture. For their test they chose alpha particles. Alpha particles are
helium atoms that have lost two electrons and have a double posi-
tive charge from the remaining two protons. In their experiment,
they directed a narrow beam of alpha particles at a very thin sheet of
gold foil. According to the existing theory, they expected the alpha
particles to pass straight through the gold.
To everyone's surprise, a small fraction of the alpha particles
deflected, or bounced off, the gold foil at very large angles. A fevv
alpha particles even bounced straight back toward the source. In
Figure 4.6, you can see an illustration of Rutherford's apparatus.
Source ofalpha
particles
Lead shield
(a)
Beam ofalpha
particles
Alpha particles
Gold foil
Fluorescentscreen
Nucleus
Atoms Ofgold foil
(b)
I-I 1.1
Based on the experimental results, Rutherford suggested a new
theory of the atom. He proposed that almost all the mass and all the
positive charge are concentrated in a small region at the center of
the atom. He called this region the nucleus. The nucleus is the cen-
tral core of an atom, composed of protons and neutrons. Because
protons and neutrons have a much greater mass than electrons,
almost all of the mass of an atom is concentrated in a tiny nucleus.
The nucleus is so dense that if it were the size of a pea its mass
would be 2.3 x 105 kg (250 tons)! In Figure 4.7 the size of an atom is
compared to that of a football stadium.
The nucleus has a positive charge, and it occupies a very small
part of the volume of an atom. What about the region of the atom
beyond the nucleus? Rutherford thought that the rest of the atom
was more or less empty space. The negatively charged electrons in
that area occupied most of the volume of the atom, but they were so
small that they did not interfere with the movement of the alpha
particles. So most of the alpha particles passed through unde-flected. Only when an alpha particle came close to the dense,positively charged nucleus was it deflected. Alpha particles that
made a "direct hit" on a gold nucleus bounced straight back.Rutherford later recollected: "It was about as credible as if you had
fired a 15-inch shell at a piece of tissue paper and it came back and
hit you."
Figure 4.7 If an atom were the
size of this stadium, then its
nucleus would be about the size
of a marble.
SafetyProper shielding shouldalways be used with radio-active emissions such as
alpha particles.
4.3 The Structure of the Nuclear Atom 91
Objective
Explain how the atomicnumber identifies an element,
concept Practice
5. How did the results of Rutherford's gold foil experiment
differ from his expectations?
6. What is the charge, positive or negative, of the nucleus of
every atom?
Recall that most atoms are composed of electrons, protons, and
neutrons. The protons and neutrons make up the small, densenucleus. The electrons surround the nucleus and occupy most of
the volume of the atom. How are atoms of one element different
from those of another element? The answer is that differencesamong elements result from differences in the numbers of protons
in their atoms. As you can see in Table 4.2, atoms of boron (B) have
five protons, atoms of carbon (C) have six protons, and fluorine (F)
atoms have nine protons.What makes an atom a hydrogen atom? Every hydrogen atom
has one proton in its nucleus. Every oxygen atom has eight protonsin its nucleus. The atomic number of an element is the number ofprotons in the nucleus of the atom of that element. Since all hydro-gen atoms have one proton, the atomic number of hydrogen is 1.
Table 4.2_
Name
Hydrogen
Helium
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
Atoms of the First Ten Elements
Composition of the nucleus
Symbol
Be
c
o
Ne
Atomicnumber Protons Neutrons*
1 1
2 2
3 3
4 4
5 5
6 6
7 7
8 8
9 9
10 10
2
4
5
6
6
7
8
10
10
Massnumber
1
4
7
9
11
12
14
16
19
20
Number ofelectrons
2
3
4
5
6
7
8
9
10
• Number of neutrons in the most abundant isotope. Isotopes are introduced in Section 46,
92 Chapter 4 Atomic Structure
Similarly, since oxygen atoms have eight protons, the atomicnumber of oxygen is 8. The atomic number identifies an element.Remember that atoms are electrically neutral. Thus the number ofprotons (positively charged particles) in the nucleus ofan atom mustequal the number of electrons (negatively charged particles) aroundits nucleus. A hydrogen atom has one electron around its nucleus,and an oxygen atom has eight electrons.
The periodic table also gives the atomic number of each ele-ment. Notice that the atomic number is a whole number writtenabove the chemical symbol of each element. The atomic numberincreases as you read across each row of the periodic table from leftto right.
Example 1 Finding Numbers of Protons and Electrons
The element nitrogen (N) is atomic number 7. How many pro-tons and electrons are in a nitrogen atom?
v otion
The atomic number equals the number of protons or thenumber of electrons in an atom. Since the atomic number is 7,a nitrogen atom has seven protons and seven electrons.
Concept Practice
7. Why is an atom electrically neutral?
B. What is the relationship between the number of protons
and the atomic number of an atom?
Practice problem
9. Use the periodic table to complete this table.
Element Symbol Atomicnumber
potassium
5
Number ofprotons
16
4.4 Atomic Number 93
Infer the number of protons,electrons, and neutrons usingthe atomic number and massnumber of an element.
94 Chapter 4 Atomic Structure
4.5 Mass NumberYou know that most of the mass of an atom is concentrated in its
nucleus and depends on the number of protons and neutrons. Table
4.2 shows that a helium atom has two protons and two neutrons
and a mass number of 4. A carbon atom with six protons and six
neutrons in its nucleus has a mass number of 12. Thus the total
number of protons and neutrons in the nucleus is the mass number
ofan atom.You can determine the composition of an atom of any element
from its atomic number and its mass number. The atom of oxygen
shown in Table 4.2 has an atomic number of eight and a massnumber of 16. Since the atomic number equals the number of pro-
tons and the number of electrons, an oxygen atom has eight protons
and eight electrons. How can you find the number of neutrons inthis atom? The mass number is 16 and is equal to the number of
protons plus the number of neutrons. Oxygen, then, has eight neu-
trons, the difference between the mass number and the atomicnumber. For any atom:
Number of neutrons = mass number — atomic number
Example 2 Determining the Composition of an Atom
How many protons, electrons, and neutrons are in thefollowing atoms?
a. Beryllium (Be)
b. Neon (Ne)
c. Sodium (Na)
Solution
Atomic number
4
10
11
Mass number
9
20
23
a. The atomic number equals the number of electrons and
the number of protons. So Be has four protons and four
electrons. To find the number of neutrons, subtract the atomic
number from the mass number. Thus, this Be atom has 9 — 4 =
5 neutrons.b. Using the same reasoning, this atom of Ne has 10 protons,
10 electrons, and 20 — 10 = 10 neutrons.
c. The atom of Na has Il protons, Il electrons, and 23 — Il =
12 neutrons.
Practice Problem
10. Complete this table.
NumberAtomic Mass ofnumber number rotons
9
14
47
55 25
Number Number Symbolof of of
neutrons electrons element
10
15
22
To represent the composition of any atom in shorthand nota-tion, you use the chemical symbol with two additional numberswritten to the left of it. The atomic number is written as a subscript(a number lowered slightly). The mass number is written as a super-
script (a number raised slightly). Look at Figure 4.8. How manyneutrons does gold have? Atoms of hydrogen with a mass number of
I may be designated hydrogen-I. Atoms of helium with a massnumber of 4 are designated helium-4.
Example 3 in Determining
an Atomthe Number of Neutrons
How many neutrons are in the following atoms?168 • 16 c. d. 382Br e. 230)b
Sonnion
Recall that the superscript is the mass number and the sub-
script is the atomic number. The mass number minus the
atomic number equals the number of neutrons.
b. 16 c. 61 d.45 e. 125
Concept Practice
11. An atom is identified as platinum-195.
a. What is the number 195 called?
b. Write the symbol for this atom using superscripts andsubscripts.
197
79
Figure 4.8 Au is the chemicalsymbol for gold. How manyelectrons does gold have?
When the composition of anatom is represented in short-hand form, subtract thebottom number from the topnumber to get the numberof neutrons.
4.5 Mass Number 95
Objective
Explain how isotopes of anelement differ
1 on o
10p•o11rP'
e IOp•i2n0
toe-
Neon-2010 protons10 neutrons10 electrons
Neon-2110 protons11 neutrons10 electrons
Neon-2210 protons12 neutrons10 electrons
Figure 4.9 Neon-20, neon-21,and neon-22 are three isotopes ofneon. How do these isotopesdiffer? How are they the same?
96 Chapter 4 Atomic Structure
Practice Problem
12. Determine the number of neutrons in each atom.
a. carbon-13 b. nitrogen-15 c. radium-226
Most of Dalton's atomic theory is still accepted today. It is now
known, however, that atoms of the same element may have different
nuclear structures. The nuclei of the atoms of a given element must
all contain the same number of protons, but the number of neu-
trons may vary. You can see in Figure 4.9 that there are threedifferent neon atoms. How do these atoms differ? Each of the neon
atoms has the same number of protons (10) and electrons (10) but a
different number of neutrons. Atoms that have the same number of
protons but different numbers of neutrons are called isotopes.
Because isotopes of an element have different numbers of neutrons,
they also have different mass numbers. Despite these differences,
isotopes are chemically alike because they have identical numbers
of protons and electrons. These subatomic particles are responsible
for the characteristic chemical behavior of each element.
Look for hydrogen in Table 4.3. There are three known isotopes
of hydrogen. Each isotope of hydrogen has one proton in thenucleus. The most common hydrogen isotope has no neutrons. Ithas a mass number of 1 and is called hydrogen-I (}H), or hydrogen.
The second isotope has one neutron and a mass number of 2. It iscalled either hydrogen-2 (21H) or deuterium. The third isotope hastwo neutrons and a mass number of 3. This isotope is hydrogen-3(fH) or tritium.
Example 4 Writing Formulas of Isotopes
Two of the isotopes of carbon are carbon-12 and carbon-13.
Give the chemical symbol for each.
Solution
The mass numbers are given in the names of the isotopes.
Carbon is atomic number 6. All atoms of carbon have sixprotons.
Carbon-12, lic Carbon- 13,
Table 4.3 Natural Percent Abundance of Stable Isotopes of Some Elements
Name Symbol
Hydrogen
Helium
Carbon
licNitrogen
Oxygen
1780
IRO
Sulfur fis
34 S16
36
16S
Chlorine
Zinc 64 Zn66 Zn67 Zn68 Zn
738zn
concept Practice
Mass(amu)
1.0078
2.0141
3.0160
3.0160
4.0026
12.000
13.003
14.003
15.000
15.995
16.995
17.999
31.972
32.971
33.967
35.967
34.969
36.966
63.929
65.926
66.927
67.925
69.925
Naturalpercent abundance
99.985
0.015
negligible
0.0001
99.9999
98.89
1.11
99.63
0.37
99.759
0.037
0.204
95.00
0.76
4.22
0.014
75.77
24.23
48.89
27.81
4.11
18.57
0.62
13. How are isotopes of the same element alike? How are they
different?
Practice Problems
14. Three isotopes of oxygen are oxygen-16, oxygen-17, and
oxygen-18. Write the chemical symbol, including the atomic
number and mass number, for each.
15. Use Table 4.3 to determine the number of protons, elec-
'Average"atomic mass
1 0079
4.0026
12.011
14.007
15.999
32.064
35.453
65.37
trons, and neutrons in each of the five isotopes of zinc.
4.6 Isotopes of the Elements 97
-7
About 1200 sitesthroughout the United Stateshave been designated as prioritycleanup sites because theythreaten human health or theenvironment.
98 Chapter 4 Atomic Structure
a.A Environmental AwarenessHazardous WastesIn the United States there are thousands of hazardous waste sites in
need of cleanup. Cleaning up waste sites is an extremely time-con-
suming and expensive process. In 1980, Congress passed theComprehensive Environmental Responsibility, Compensation and
Liability Act (CERCLA) to address the problem of hazardous wastecleanup. Also known as Superfund, this law authorizes theEnvironmental Protection Agency to clean up abandoned haz-ardous waste sites.
When a potential site is identified, it must first be investigated
to see if it poses a threat to human health or to the environment. If
so, the site is placed on a national priority cleanup list. By 1990, one
decade after the passage of Superfund, approximately 1200 sites
had been tabbed for the priority list. About 50 of the sites on the pri-
ority list had been cleaned up. Individual states are cleaning up an
additional 1000 sites.
Cleaning up a hazardous waste site is a complex project. One
project may bring together experts from the fields of chemistry,biology, hydrogeology, engineering, medicine, toxicology, law, and
politics. The experts study the site and decide on the best cleanupapproach to take. Traditional cleanup methods include incinera-
tion, chemical extraction, physical containment, and ground water"pump and treat."
Bioremediation is a method currently being developed todestroy toxic wastes. It involves the use of bacteria that naturally
degrade hazardous materials. The cleanup site is sprayed with spe-cial fertilizers that encourage bacterial growth. The degradationprocess can be accelerated up to one million times the rate ofnormal degradation. Bioremediation has been effective in treatingoil spills and city sewage. This method may also become a usefultool in cleaning up toxic wastes. With a blend of old and new tech-nology, plus plenty of time and money, contaminated land and
water can be restored to their pristine state.
Think About It
16. Apply An old adage says "An ounce of prevention is worth a
pound of cure." How would you apply this saying to the problem of
hazardous waste disposal?
17. Compare Of the cleanup methods listed, which do you think
would be most suited to the cleanup of contaminated water? To the
cleanup of contaminated land?
4.7 Atomic MassThe mass of even the largest single atom is much too small to bemeasured individually on a balance. A glance back at Table 4.1shows that the actual mass of a proton or a neutron is very small:1.67 x 10 24 g. Even compared with this small mass, the mass of anelectron is negligible: 9.11 x 10-28 g. Since the 1920s it has beenpossible to determine the mass of an individual atom by using amass spectrometer. The mass of a fluorine atom was found to be3.155 x 10-23 g and the mass of an arsenic atom is 1.244 x 10-22 g.
The masses of individual atoms are useful information, butthese values are inconvenient and impractical to work with. Instead,it is more useful to compare the relative masses of atoms using anisotope of carbon, carbon-12, as a basis. This isotope of carbon wasassigned a mass of exactly 12.00000 amu. An qtomic mass unit(amu) is defined as one-twelfth the mass of a carbon-12 atom. Usingthese units, a helium-4 atom, with a mass of 4 amu, has about one-third the mass of a carbon-12 atom. How many carbon-12 atomswould have about the same mass as a nickel-60 atom?
You know that a carbon-12 atom has six protons and six neu-trons in its nucleus and its mass is set as 12.00000 amu. Thereforethe mass of a single proton or a single neutron is about 1 amu.Because the mass of any single atom depends on the number ofprotons and neutrons in its nucleus, you might predict that theatomic mass of an atom should be a whole number. As you can seefrom the periodic table, the atomic masses of sodium, phosphorus,and gold are 22.990 amu, 30.974 amu, and 196.97 amu, respectively.
Each of these masses is close to a whole number. However, theatomic mass of chlorine (Cl) is 35.453 amu. How can this atomicmass be explained?
Consider the three isotopes of hydrogen discussed in the last
section. According to the table, almost all naturally occurringhydrogen (more than 99.98%) is hydrogen-I. The other two isotopes
are present in only trace amounts. In nature, most elements occur as
a mixture of two or more isotopes. Each isotope of an element has a
fixed mass and a natural percent abundance. Notice that the atomic
mass of hydrogen in the periodic table or in Table 4.3 (1.0079 amu)
is very close to the mass of hydrogen-I (1.0078 amu). The slight dif-
ference takes into account the larger masses and lower amounts of
the other two isotopes of hydrogen.
Now consider the two isotopes of chlorine: chlorine-35 andchlorine-37. The simple arithmetic average of the masses of thesetwo isotopes is 36.9674 amu [(34.9689 amu + 36.9659 amu)/21. The
atomic mass of chlorine would average 36.9674 amu only if the two
types of chlorine atoms were present in nature in equal amounts. In
reality, approximately 75% of all chlorine atoms in nature have a
Objective
Explain. using the concept ofisotopes, why the atomtc
masses ot elements are notwhole numbers.
Philosophy of SciencePhilosophers of science areprimarily concerned With thecritical analysis of scientificconcepts and the way inwhich these concepts areexpressed. These philoso-phers have analyzed suchconcepts as "number,""space," "force," and "livingorganism." For example, isthe modern scientific methodthe best way to examine andexplain the natural world? Oris there a better way, as yetunknown? Suppose thephilosophers of sciencesuggested an improved sci-entific method. Using the newmethod, scientists might gainstartling new insights andmake advances in scienceand technology. For thisreason, philosophical ques-tions and answers about themethods and concepts ofscience could be an impor-tant aid to scientific progress.
4.7 Atomic Mass 99
Objective
Calculate the average atomicmass of an element fromisotope data.
mass of approximately 35 amu (chlorine-35) and 25% have a mass
of approximately 37 amu (chlorine-37). Because there is more of the
chlorine-35 isotope, the atomic mass should be closer to 35 amuthan to 37 amu. The atomic mass of chlorine is 35.453 amu, which is
a weighted average mass of these two isotopes. The atomic mass of
an element is a weighted average mass of the atoms in a naturallyoccurring sample of the element. A weighted average mass reflects
both the mass and the relative abundance of the isotopes as they
occur in nature. The next section discusses the actual calculation of
the atomic mass of an element from isotope data.
Example 5 Finding the Isotope of Greatest Abundance
Copper has two isotopes: copper-63 and copper-65. Given that
the atomic mass of copper from the periodic table is 63.546
amu, which of the isotopes of copper is most abundant?
Solution
The atomic mass of 63.546 amu is closer to 63 than to 65, so
most of the copper atoms must be copper-63.
concept Practice
18. What data must you have about the isotopes of an elementto be able to calculate the atomic mass of the element?
19. There are three isotopes of silicon with mass numbers of
28, 29, and 30. The atomic mass of silicon is 28.086 amu.Comment on the relative abundance of these three isotopes.
4.8 calculating the Atomic Massof an ElementYour grade in a class may be calculated as a weighted average. Thatis, some exams may have more "weight," or importance, than others.For example, your teacher might count the grade on a chapter testtwo times, but count a grade on a quiz or lab only once when calcu-lating your grade average. In this way, your "average" reflects theextra weight or value of the chapter test. In the last section, you sawthat the atomic mass of an element is a weighted average of themasses of its isotopes. How is atomic mass calculated?
100 Chapter 4 Atomic Structure
Because the atomic mass must reflect both the masses and therelative natural abundances of the isotopes, you must know:
• The number of stable isotopes of that element
• The mass of each isotope
• The natural percent abundance of each isotope
You can look up both the mass and relative abundance values instandard chemistry reference books. Recall that Table 4.3 gives themass and natural percent abundance and "average" atomic mass fora few elements.
In Figure 4.11 the atomic mass of chlorine is estimated much asyour weighted grade average would be calculated. The numberobtained in Figure 4.11 is not 35.453 amu because the mass num-bers used for the estimation only approximate the actual masses ofthe isotopes and the ratio of the isotopes' abundances is not exactly3 to l. Example 6 shows a more accurate procedure for calculatingatomic masses.
Example 6 Calculating Atomic Masses
Element X has two natural isotopes. The isotope with mass10.012 amu has a relative abundance of 19.91%. The isotope,with mass 11.009 has a relative abundance of 80.09%. Calcu-
late the atomic mass of this element and name it.
Find the mass that each isotope contributes to the weightedaverage by multiplying the mass by its relative abundance.
Ratio of chlorjne atoms in naturalabundance: three 357Cl to one 31;Cl1
17p
180
17P
O18n
17p
o18
Total number ofprotons in three
3517Cl atoms andone 3717Cl atom(17+17+17+17)
o
Total number ofneutrons in three17Cl atoms andone 3ßCl atom(18 +18 +18 +20)
Then add the products.
lox 10.012 amu x 0.1991
llx 11.009 amux 0.8009
Total
1.993 amu
8.817 amu
10.810 amu
Element X is boron, atomic number 5.
Practice Problem
20. The element copper contains the naturally occurring
isotopes 29Cu and 92%Cu. The relative abundances and atomic
masses are 69.2% (mass = 62.93 amu) and 30.8% (mass =
64.93 amu), respectively. Calculate the average atomic mass
68 + 74= 35.5 amu4
Average massof one atom
Figure 4.11 The ratio Ofchlorine-35 (75% abundance)to chlorine-37 (25% abundance)is about 3 to 1. Therefore themass of three chlorine-35 atomsis averaged with the mass of onechlorine-37 atom.
of copper.
4.8 Calculating the Atomic Mass of an Element 101
Separatedcomponents
of mixture
RecorderColumn
packed with ,absorbent
Sampleinjector
Heatedmetal block
Regulatorvalve
Directionof gas
flow
Detector
To massspectrometer
Carrier
tank
The gas chromato-graph (GC) IS used to separatecomponents ot mixtures.
The mass spec-trometer (MS) is used to find themasses of the components of thegaseous mixture. When themasses of the gaseous compo-
nents are determined, thecomponents can be accuratelyidentified.
40B Science, Technctcgy, and sctietyDrug TestingTests used to identify traces of drugs of abuse in the body must be
extremely accurate. A false-positive result could ruin a career. A
false-negative result could endanger lives. The best method
currently available to test for drugs of abuse is to use gas chro-
matography combined with mass spectrometry, or GC/MS. The gas
chromatograph separates a chemical mixture into its individual
components. First, the mixture is vaporized through a separation
column and carried by another gas. The time required for a gaseous
component to pass through the column, called retention time, is
different for each compound. Retention time helps to identify the
compounds in the mixture. Alone, however, the GC method is not
very reliable. The component gases are swept from the GC into a
mass spectrometer. The MS splits the different kinds of gaseous
molecules into ions of different masses and exposes them to a mag-
netic field. The degree of deflection of an ion in a magnetic field is
related to its mass. The identity of a component can be determined
by comparing patterns of ionic deflection to the MS "deflection" of
a known substance.Used together, the GC/MS method of drug testing is very reli-
able—nearly 100%. But to some critics, almost 100% is not good
enough. For example, poppy seeds contain small amounts of a com-
pound that is extremely similar to opium and heroin. Based on aGC/MS test, it is remotely possible for a person who has just eaten a
roll with poppy seeds to be mistaken for a heroin user! Thus chemists
are busy devising new drug tests and working to improve the old ones.
Electrical
field
Gas
inlet
Electrongun
Think About It
Detector
Slit
Separated beamsof different
masses
electromagnets
21. Criticize Do you think a reliability of 95% in a drug test is ade-quate for an employer to make decisions about whether a personshould be hired? Explain your answer.
102 Chapter 4 Atomic Structure
Chapter a ReviewAtomic structure
Key Terms
atom 4.1atomic mass 4.7atomic mass unit 4.7atomic number 4.4cathode ray
2 7 _ c)
Z R 3 S
electron 4.2isotope 4.6mass number 4.5neutron 4.2nucleus 4.3
Dalton's atomic theory 4.1 proton 4.2
Chapter Summary
Atoms are the basic building blocks of matter.Each element is composed of atoms. The atoms
of a given element are different from the atoms of
all other elements.Atoms are exceedingly small. Dalton theo-
rized that atoms were indivisible, but thediscovery of the electron changed this theory.Besides negatively charged electrons, atoms con-
tain positively charged protons and electrically
neutral neutrons. The proton has a mass nearly2000 times the mass of an electron. A proton and
a neutron are nearly identical in mass.
The nucleus of the atom is composed of pro-
tons and neutrons. The nucleus contains most of
the mass of the atom in a very small volume. The
electrons surround the nucleus and occupy most
of the volume of the atom.
The number of protons in the nucleus of the
atom is the atomic number of that element.Atoms are electrically neutral. Thus an atom has
the same number of protons and electrons. The
sum of the protons and neutrons is the massnumber. The atoms of a given element all con-
tain the same number of protons, but thenumber of neutrons may vary. Atoms with thesame number of protons but different numbers
of neutrons are isotopes.
The atomic mass of an element is expressedin atomic mass units (amu). An atom of any ele-
ment has an atomic mass that is approximately a
whole number. This is because protons and neu-
trons each have a mass of about I amu. The
atomic mass in the periodic table is a weighted
average of all the naturally occurring isotopes of
that element. For this reason, the atomic mass of
most elements is generally not a whole number.
Practice Questions and Problems
22. In your own words state the main ideas ofDalton's atomic theory. 4. I
23. Would you expect two electrons to attract or
repel one another? 4.224. What are the charges and relative masses of
the three subatomic particles that are of most
interest to chemists? 4.225. What did Rutherford's gold foil experiment tell
us about the structure of the atom? 4.326. Describe the composition of the nucleus of
the atom. 4.327. What does the atomic number of each atom
represent? 4.428. What is meant by the statement "Atoms are
electrically neutral"? 4.4
29. What is the difference between mass number
and atomic number? 4.530. Complete this table. 4 5
Number Mass Number Atomic NumberElement protons number electrons number neutrons
Si 15
1 2
50 24
88 38
31. Name two ways in which isotopes of an ele-ment differ. 4.6
32. fist the number of protons, neutrons, andelectrons in each of the following atoms. 4.6
a. {Alb 42tca
Chapter 4 Review 103
33. What is an atomic mass unit? 4.734. What is the atomic mass of an element? 4.735. Uranium has three isotopes with the following
percent abundances: 2}2tJ (0.0058%),
(0.71%), and (99.23%). What do youexpect the atomic mass of uranium to be in
whole numbers? Why? 4.736. What information about an element's isotopes
is needed to calculate that element's atomic
mass? 4.8
Mastery Questions and Problems
37. Make a concept map using atom as the main
concept. Use the chapter key terms and the
terms negative, positive, and neutral in your
map.38. Explain why the atomic masses of most ele-
ments are not whole numbers.
39. Compare the relative size and relative density
of an atom to its nucleus.
40. How can there be more than 1000 different
atoms when there are only about 100 different
elements?
41. Imagine you are standing on top of a boron-11
nucleus. Describe the numbers and kinds of
subatomic particles you would see when you
look down into the nucleus and those you
would see when you look out from the
nucleus.
42. What parts of Dalton's atomic theory no longer
agree with our current picture of the atom?
43. The four isotopes of lead are shown below,
each with its percent by mass abundance and
the composition of its nucleus. Using these
data, calculate the atomic mass of lead.
26.26% 20.82% 51.55%
44. Dalton's atomic theory was not correct inevery detail. Should this be taken as a criti-
cism of Dalton as a scientist? Explain why or
why not.
104 Chapter 4 Atomic Structure
45. Why are atoms considered the "basic building
blocks" of matter even though smaller parti-
cles, such as protons and electrons, exist?
46. The following table shows some data collected
by Rutherford and his colleagues during their
gold foil experiment.
a. What percent of the alpha particle deflec-
tions were 50 or less?
b. What percent were 150 or less?
c. What percent were 600 or greater? C; O
Angle of deflection(degrees)
5
10
15
30
45
60
75
105
Number of deflections
8 289 000
502 570
120 570
7 800
1 435
477
211
198
Critical Thinking Questions
47. Choose the term that best completes thesecond relationship.a. female:male
(l) atom(2) neutron
b. cow:horse(l) proton(2) nucleus
c. atom:proton(l) school(2) nucleus
proton:(3) electron(4) quarkneutron:
(3) atom(4) quarkhouse:(3) planet(4) brick
48. How could you modify Rutherford's experi-
mental procedure to determine the relativesizes of different nuclei?
49. Criticize the statement: "You can't see atoms."
50. Rutherford's atomic theory proposed a dense
nucleus surrounded by very small electrons.
This implies that atoms are composed mainly
of empty space. If all matter is mainly empty
space, why is it impossible to walk through
walls or pass your hand through your desk?
51. What happens when new experimental resultscannot be explained by the existing theory?Base your answer on the scientific method.
52. The goal of environmental cleanup under theSuperfund is to reclaim land, air, and waterfrom chemical pollution. If efforts underSuperfund continue at the present rate, howlong would it take for the 1200 sites to becleaned up? In light of this time estimate, write
a paragraph supporting the current pace or anincreased pace for these cleanup efforts.
53. Of every 20 drug nonabusers tested, howmany might have a false-positive test result ifthe drug test is 95% reliable? Would it be betterto take more than one drug test? Explain youranswer in terms of the accuracy and precisionof the drug test.
Cumulative Review
54. Oxygen and hydrogen react explosively toform water. In one reaction, 6 g of hydrogencombined with oxygen to form 54 g of water.How much oxygen was used?
55. How many significant figures are in each ofthese measurements?a. 4.607 mg c. 0.00150 ml.b. 4.35 x 104 km d. 60.09 kg
56. Round each of the measurements in Problem55 to two significant figures.
57. The law of conservation of mass was intro-duced in Chapter 1. Use Dalton's atomictheory to explain this law.
58. An aquarium measures 55.0 cm x 1.10 m x
80.0 cm. How many cm3 of water will thisaquarium hold?
59. What is the mass of 5.42 cm3 of platinum? The
density of platinum is 22.5 g/cm3.
60. Classify each of the following as an element,
compound, or mixture.a. sulfur c. newspaper e. cardboardb. salad oil d. orange f. apple juice
Challenge Questions and Problems
61. Lithium has two naturally occurring isotopes.
Lithium-6 has an atomic mass of 6.015 amu;
lithium-7 has an atomic mass of 7.016 amu.The atomic mass of lithium is 6.941 amu.What is the percentage of lithium-7 in nature?
62. When the masses of the particles that make
up an atom are added together, the sum isalways larger than the actual mass of theatom. The "missing" mass, called the mass
defect, represents the matter converted intoenergy when the nucleus was formed from its
component protons and neutrons. Calculate
the mass defect of a chlorine-35 atom by using
the data in Table 4.1. The actual mass of achlorine-35 atom is 5.81 x 10-23 g.
Connections Questions
63. What was John Dalton's vocation?
64. Is it correct to say that atoms are composed ofonly three subatomic particles? Explain.
65. How might a philosopher of science aid sci-entific progress?
Write About Chemistry
66. Imagine that you are a newspaper journalist inthe early 1900s. Write a 250-word account ofthe discovery of the neutron for your paper.Hint: Journalists try to answer five questions:Who? What? When? Where? Why?
67. Write an imaginary TV interview in whichJohn Dalton defends his atomic theory to you,the interviewer.
Readings and References
Asimov, Isaac. Atom: Journey Across the Sub-atomic Cosmos. New York: Dutton, 1991.
Berger, Melvin. Atoms, Molecules, and Quarks.New York: Putnam, 1986.
Chapter 4 Review 105
