Download Chapter 24: The Nucleus

24. The Nucleus
The last several chapters have shown that an understanding of the behavior of atomic electrons can lead to
an understanding of the regularities associated with
chemical reactions as well as the structure and function
of materials.
However, in all of the chemistry that we have studied, even in the most violent of chemical reactions, the
nucleus remains untouched and unaffected at the center
of the atoms. Nevertheless, most of the mass-energy of
the atom is locked in the nucleus, and an understanding
of its structure leads us to practical applications which
have become very important to the well-being of human
beings. Nuclear weapons, nuclear power plants, cancer
treatments, and radioactive dating methods have all
sprung from the study of the nucleus.
1.
The number of protons in a nucleus is the same
as the atomic number of the atom. This determines the number of electrons in the neutral
atom and, thus, the chemical properties of the
atom. All nuclei of a particular chemical element have the same number of protons.
2.
Most nuclei contain one or more neutrons in
addition to the protons. The neutrons add mass
to the nucleus, but not electric charge.
3.
Both protons and neutrons are called nucleons.
4.
The total number of nucleons (protons and neutrons) in a nucleus is called its mass number.
Protons and Neutrons
5.
Different isotopes of a particular element have
the same number of protons in each nucleus,
but different numbers of neutrons. For example, all oxygen nuclei have eight protons, but
some have eight neutrons, others have nine
neutrons, and still others have ten neutrons.
The mass numbers of these isotopes are 16, 17,
and 18, respectively.
6.
The different isotopes of a particular element
are usually designated by adding the mass
number as a superscript to the element’s symbol. The atomic number may also be designated by a subscript. For example, the oxygen
isotopes described above would be designated
as 168O, 178O, and 188O.
Chapter 17 explained that each neutral atom is
characterized by a number of electrons that occupy
orbitals about the atomic nucleus, and that the number
of protons is the atomic number of the atom. This determines the chemical element that the atom represents
and, hence, its chemical properties.
The positive charge in each atom comes from the
protons in the atomic nucleus. The amount of positive
charge is balanced by the negative charge of the electrons
so that the atom as a whole is electrically neutral. Each
nucleus may also contain neutrons—particles that are
electrically neutral, and have about the same mass as protons.
The nuclei corresponding to a particular element,
such as fluorine, may all have the same mass. However,
most elements can have more than one kind of nucleus;
the different nuclei of an element are called isotopes,
some of which are shown in Table 24.1. There are three
kinds of hydrogen nuclei. Most helium nuclei have the
mass of four protons, but some have only that of three.
There are three kinds of carbon nuclei, two kinds of
copper, five of zinc, eight of tin, and nine of xenon. The
nuclei of different isotopes of an element have the same
number of protons, but different numbers of neutrons.
The structure of atomic nuclei can be summarized
as follows:
Naturally occurring elements usually contain several isotopes, all of which have the same chemical properties. The chemical atomic mass of an element is the
average atomic mass of the isotopes that make up the
element. In the case of copper, which has two isotopes,
69 percent of the copper atoms have a mass number of
63, and 31 percent have a mass number of 65. All have
29 protons, but some have 34 neutrons and others have
36. The average mass of copper atoms, taking the relative abundances of the two isotopes into account, is
63.54. This is the number published as the “atomic
mass” of copper, since both kinds of copper atoms are
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Table 24.1. Some isotopes of the elements.
ATOMIC
NUMBER
1
1
1
ISOTOPE
NUMBER OF
PROTONS
1
H
1H
3
1H
1
2
NUMBER OF
NEUTRONS
MASS
NUMBER
1
1
1
0
1
2
1
2
3
2
2
1
2
3
4
2
2
3
6
6
6
12
C
6C
14
6C
6
6
6
6
7
8
12
13
14
8
8
8
16
O
8O
18
8O
8
8
8
8
9
10
16
17
18
9
19
F
9
10
19
Cu
29Cu
29
29
34
36
63
65
92
92
143
146
235
238
He
2He
2
4
6
13
8
17
9
29
29
63
92
92
235
29
65
U
92U
92
238
present in all chemical reactions.
nuclei and atoms is that nuclei have more ways to
release energy, some of which are discussed below.
Radioactivity
Alpha Decay
Most of the atomic nuclei in nature are stable; they
do not change if left to themselves. They are not even
affected by violent chemical reactions such as explosions or combustion.
These reactions involve
rearrangements of atomic electrons, but not changes in
the nuclei themselves.
However, some naturally occurring nuclei spontaneously undergo changes that result in rearrangement of
the nuclear constituents and the release of significant
amounts of energy. Such nuclei are said to be radioactive.
The situation seems to be something like that which
occurs when atomic electrons are in high-energy states.
Such excited electrons can return to the lower energy
states only if their excess energy can be released, perhaps by creating and emitting a photon. Radioactive
nuclei are also in states with more energy than necessary. They can become more stable if the excess energy can be released. However, the difference between
The nucleus may emit a fast, massive particle that
contains two protons and two neutrons (Fig. 24.1). When
first discovered these particles were called “alpha rays,”
but they have since been found to be identical to the
nuclei of 4He atoms. Even today, such nuclei when produced in nuclear processes are called alpha particles.
Rutherford used these naturally occurring alpha
particles as the “bullets” in his early experiments on
atomic nuclei discussed in Chapter 15. Since the nucleus loses two protons in alpha decay, the resulting nucleus (“daughter”) has a lower atomic number than before
and thus belongs to a different chemical element. Its
mass number is reduced by four. An example is the
radioactive decay of radium.
226
Ra →
88
222
Rn + 42He .
86
Note that both charge and mass number are conserved
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Figure 24.1. Alpha decay.
Figure 24.2. Beta decay.
in this transformation.
They interact weakly and only with nuclear particles.
As a result they are able to penetrate large thicknesses
of material, such as the entire diameter of the earth, with
only a small probability of interacting with anything.
Neutrinos are created and emitted in beta decay and
carry away some of the excess energy of the decaying
nucleus.
Beta Decay
The nucleus may emit a fast electron (Fig. 24.2).
Since there is reason to believe that electrons cannot be
confined in the nucleus (using the Uncertainty
Principle, for example), a neutron within the nucleus
appears to create and immediately emit the electron,
much as orbital electrons create and emit photons as an
energy release mechanism. Neutrons isolated outside of
atomic nuclei always decay by beta emission after a
short time. The neutrons become protons, emitting
energy and negative charge in the form of fast electrons
Gamma Decay
Gamma decay is most like the emission of light by
atomic electrons (Fig. 24.3). Alpha and beta decay usually leave the particles of the “daughter” nucleus in
excited states. The daughter nucleus can move to
lower-energy states by emitting a photon. However,
nuclear states usually involve greater energy changes
than electron states in atoms, so the resulting photons
from nuclei have higher energy than those emitted by
atoms. These high-energy photons are called gamma
rays.
Neither the atomic number (number of protons) nor
the mass number (total number of nucleons) of nuclei
changes during gamma decay, although mass does
change because energy is released. Gamma decay usually follows all the other radioactive decays, because the
residual nuclei are almost always left in an excited condition.
n → 11p + 0–1e .
1
0
This is the basic beta-decay reaction. It sometimes
occurs with neutrons inside nuclei. The resulting nucleus will have one fewer neutron and one more proton
than before. Again, it will belong to a new chemical
element. An important example is the beta decay of carbon-14.
14
C → 147N + 0–1e .
6
Once again note that mass number and electric charge
are both conserved.
Electron Capture
Subsequent research has shown that another particle, called a neutrino, is also emitted with each electron
in beta decay. Neutrinos have little or no rest mass, no
electric charge, and travel at or near the speed of light.
Sometimes the nucleons could have a lower energy
if one of the protons could become a neutron (Fig.
24.4). One mechanism by which this can occur is for
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Figure 24.3. Gamma decay.
the nucleus to “capture” one of the orbiting atomic electrons, combining the electron with a proton to make a
neutron. The excess energy is emitted in the form of a
neutrino and, sometimes, one or more gamma rays.
The result of electron capture is a decrease in the
atomic number, since the nucleus now has less charge
than before. As with beta decay, the nuclear mass number does not change; there are the same total number of
nucleons as before.
An important example of electron capture occurs in
the decay of potassium-40.
40
K + 0–1e →
19
40
11
C →
6
11
B + 0+1e + 00 neutrino .
5
Positrons, which are produced by naturally occurring gamma rays in the atmosphere and in matter, have
an interesting history. As they gradually slow down,
they transfer their kinetic energy to atomic electrons and
cause ionization. When they are slow enough, they
attract an electron, and the two of them form a little
atom called positronium. (However, it is a strange kind
of atom because the electron and positron have the same
mass. Positronium is an atom without a nucleus.) After
a short time the positron and electron annihilate each
other, emitting their total mass-energy in the form of
two gamma rays.
Ar + 00neutrino .
18
Note once again that total mass number and electric
charge are both conserved in this reaction.
Fission
Positron Decay
Some nuclei, usually the heaviest ones, have so
much excess energy that they break apart into two large
fragments in a process called spontaneous nuclear fission (Figure 24.6). The products of such fission are
always neutron rich—they have too many neutrons—
and are always radioactive. They begin emitting energy, usually by beta emission. In addition, some of the
neutrons of the original fissioning nucleus are not
included in either of the major fragments. These
become free neutrons, which we will study in the next
chapter. Some nuclei that do not spontaneously decay
by fission can be made unstable by absorbing a neutron.
Important examples of such induced nuclear fission is
Nature provides one other mechanism to convert
protons in nuclei into neutrons. If the available energy
is large enough, nuclei will sometimes emit a particle
that has all the properties of an electron except that it
carries a positive rather than a negative charge. Such a
particle is called a positron.
Positron decay (Fig. 24.5) is like beta decay in
every way, except that the emitted particle has positive
charge and the resulting nucleus has an atomic number
that is one lower, rather than one higher, than before the
decay took place. For example,
Figure 24.4. Electron capture. The nucleus “captures” an orbital electron, changing one proton to a neutron.
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Figure 24.5. Positron decay and annihilation. A proton inside a nucleus becomes a neutron by emitting a positive electron. The positron later combines with a normal electron and annihilates.
Figure 24.6. Spontaneous fission.
235
U + 10n →
92
90
then become ionized (Fig. 24.7). All the effects caused
by radioactivity can be traced to either the ionized
atoms or the free electrons that are produced in this way.
(Gamma rays also cause ionization through the highenergy version of the “photoelectric” removal of electrons from atoms and, indirectly, through the production
of high-energy positron-electron pairs.) Radioactive
emissions are sometimes called ionizing radiation
because of the ionization they cause.
The oldest practical application of radioactivity is
in the making of radium-dial watches. Radioactive
material is mixed with a luminescent powder and painted on watch dials. The charged particles released by the
radioactivity separate electrons from atoms in the powder. As the electrons return to their equilibrium states,
they emit light.
Another early application was in the treatment of
cancer (Fig. 24.8). The radioactive emissions cause
ionization in biological materials as well as in nonliving
substances. Such ionization disrupts a cell’s normal
Sr + 14454Xe + 10n + 10n
38
or
U + 10n → 8838Sr + 14654Xe + 10n + 10n .
Exactly how the unstable nucleus breaks up is somewhat a matter of chance.
235
92
Application of Radioactive Materials
Radioactivity represents significant energy release
from the nuclei of atoms. Not surprisingly, such energy
has important effects in nature and is used in several
important devices.
The energy released by radioactive processes
appears as the kinetic energy of the emerging charged
particles. These transfer their kinetic energy to the matter through which they pass by interacting via the electrical interaction. Most often these fast, charged particles interact with electrons in matter, dislodging them
from the atoms to which they are attached. These atoms
Figure 24.7. Radioactive emissions transfer energy to matter by causing ionization. Which of the fundamental interactions is responsible?
227
function. Rapidly reproducing cells like cancer cells
seem particularly susceptible to this kind of damage. If
there is enough disruption, individual cells lose their
ability to function and die. This possibility makes ionizing radiation one of the most important weapons
against certain diseases, particularly cancer. Radiation
treatment alone has changed cancer of the uterine cervix
from one of the principal causes of death in women to
one of the most curable of all cancers. On the other
hand, ionizing radiation can also cause undesirable biological effects.
space probes use radioactive power cells.
Finally, radioactivity in the materials of which the
earth is composed provides the energy that keeps the interior of the earth at a higher temperature than its surface. If
there were no such source of energy, the earth would have
cooled long ago to a uniform temperature. Processes that
occur on the surface of the earth would be different if this
significant energy source were not operating.
Radioactive Half-life
The decay of radioactive nuclei is a statistical
process. The decays are governed by waves of probability, just as atomic processes are governed by orbitals
of probability. Predicting the instant that a particular
nucleus will decay is impossible. It may wait several
thousand years in its excited state, or it may decay in the
next instant. However, if there is a large collection of
similar nuclei the average behavior of the group can be
predicted and measured with considerable accuracy.
Some of the nuclei will decay almost immediately, others will decay after a short time, and still others will
wait a long time before their radioactive decay.
One way to describe the decay of a particular sample of radioactive nuclei is to specify its half-life, the
time required for half of the nuclei to decay. The halflife is a characteristic of particular species of radioactive
nuclei and varies from a fraction of a second to many
billions of years for different species of nuclei. For
example, the half-life of a sample of carbon-14 nuclei is
5,730 years, whereas that of a sample of potassium-40
nuclei is 1.3 billion years.
The statistical nature of radioactive decay has an
interesting consequence we will need to know about.
Suppose we have a certain sample of radioactive material and measure its decay as time passes (Fig. 24.9).
After a time equal to the half-life of the material, half of
the original nuclei would have decayed. At the end of a
second half-life, half of the remaining nuclei would
have decayed. In total, three-fourths of the original
material would have decayed, leaving one-fourth as it
was at the beginning. During a third half-life, half of
these would decay leaving one-eighth in the original
Figure 24.8. Gamma radiation from radioactive materials can kill cancer cells inside the body.
In more recent years radioactivity has been used to
run small power cells in applications that require small
amounts of energy over a long time, and for which it is
inconvenient to change batteries. For example,
implanted heart-pacemakers and some applications in
Figure 24.9. The random decay of radioactive nuclei. Each frame represents the passage of one half-life.
228
form. As each half-life passes, one-half of the material
present at the beginning of the interval decays.
For example, suppose that we start with 1 billion
atoms of carbon-14. (This is a small number for any
real sample.) After 5,730 years (the half-life of carbon14), 500 million atoms would remain. After the next
5,730 years, 250 million would survive; 125 million
would be around after a third 5,730 years. The decay
would continue in this way, the number of survivors
being halved every 5,730 years, as long as significant
numbers of nuclei remain in the sample. (When the
numbers become small, the laws of probability are no
longer adequate to give an accurate prediction of when
the last few nuclei will decay.)
The half-lives of radioactive nuclei depend on
processes that take place inside nuclei themselves, but
seem not to depend on ordinary processes in which the
outer atom might be involved. Events such as chemical
reactions and ambient physical conditions such as temperature and pressure do not alter the decay of unstable
atomic nuclei.
argon is trapped in the rock as a compound or mineral.
Thus, when it is cooled to solid form, we begin with a
rock that is free of argon. If we subsequently analyze a
rock containing potassium-40, the amount of argon-40
reveals the number of potassium-40 nuclei that have
decayed since the rock solidified. This, together with a
measurement of the number of potassium-40 nuclei that
remain, allows a calculation of the number of half-lives
that have elapsed. The method assumes that the
amounts of potassium-40 and argon-40 found remaining in the rock are related by radioactive decay. It also
assumes that once the original rock has cooled and
solidified, it is not again melted or subjected to nearmelting temperatures that would drive off accumulated
argon. The method is limited to measuring the ages of
rocks that had a molten origin at least some tens of millions of years ago or more.
There are several radioactive materials in the
earth’s crust that permit this same kind of calculation.
In each case it is possible to estimate or measure the
number of nuclei that have decayed and the number that
remain. The fraction of nuclei that have decayed
reveals the number of half-lives that have elapsed and
this, in turn, allows a calculation of the time interval
since the formation of the material. All of these calculations indicate that the earth is about 4.6 billion years
old.
Another important isotope used for dating is carbon-14. Unlike the radioactive materials in the earth’s
crust, carbon-14 is continuously being formed in the
earth’s atmosphere as cosmic rays (mostly protons
ejected by the sun) bombard atmospheric nitrogen
atoms. After their formation, carbon-14 atoms combine
with atmospheric oxygen to form carbon dioxide. They
may then become part of living material through the
normal carbon cycle, being incorporated by plants into
biologically important materials.
Most carbon in living things is not carbon-14 at all,
but rather carbon-12. Unlike carbon-14, carbon-12 is a
stable isotope of carbon and does not decay. For every
carbon-14 atom in the atmosphere, there are about 1012
carbon-12 atoms. The activity of the sun keeps this
fractional ratio of carbon-14 to carbon-12 constant by
generating new carbon-14 to replace that which decays.
The method assumes that the sun has produced the carbon-14 at about the same rate for the past 70,000 years,
i.e., the sun has shone with about the same intensity
over that short portion of the sun’s lifetime. Since all
living things are continuously exchanging their carbon
with the atmosphere, the fraction of radioactive carbon14 to stable carbon-12 atoms in living plants and animals also remains relatively constant.
However, when an organism dies it no longer
exchanges carbon with the atmosphere. The carbon-14
is no longer replenished as it decays, and so the fraction
Radioactive Dating
The rates at which radioactive materials decay provide a set of clocks, which can be used to estimate time
intervals under certain circumstances. For example, the
age of the earth and its materials has been debated for
hundreds of years. Radioactive dating of the earth’s
materials has finally given some reliable data from
which such estimates can be made.
The first observation is that the earth’s materials cannot be infinitely old. There are many radioactive isotopes
present in the earth’s crust. In fact, all the elements with
atomic numbers greater than 83 (bismuth) are radioactive. If the earth were infinitely old, these would all have
decayed; yet many of them are still present.
The second important observation is that the earth
is probably more than several million years old.
Radioactive materials in the earth’s crust all have halflives exceeding about 1 billion years. Several other isotopes, with half-lives in the range of a few million years,
are not present. It is argued that many of these must
have been formed at the same time as other earth materials, but they have all decayed since that time. Thus,
the minimum age of the earth is several million years
and the maximum age is several billion years.
Some radioactive isotopes permit a more precise
estimate. Potassium-40, for example, has a half-life of
1.3 billion years. When it decays by electron capture,
the product is argon, which is normally a gas. When
potassium decays inside a rock, the argon atoms are
trapped. The method is applied to rocks which begin in
the hot, molten state. The high temperatures boil off
any existing gases from the molten rock, including
argon. Since argon is a noble gas, we also know that no
229
of undecayed but unstable carbon-14 nuclei relative to
the stable carbon-12 nuclei decreases at a predictable
rate. Measurement of this slowly changing fraction permits an estimate of the time that has elapsed since the
death of the organism.
The 5,730-year half-life of carbon-14 limits the
maximum time interval for which this method is useful
to about 70,000 years (about 12 half-lives). By this time
the carbon-14 is down to about 1/4096 of its original
concentration, and the uncertainties in the resulting time
estimates increase.
4.
C. GLOSSARY
1. Alpha Decay: A mode of radioactive decay in
which a cluster of two protons and two neutrons
(alpha particle) is emitted.
2. Atomic Number: See Chapter 17.
3. Beta Decay: A mode of radioactive decay in which
a high-energy electron (beta particle) and a neutrino (technically, an antineutrino) are emitted.
4. Electron Capture: A mode of radioactive decay in
which an orbital electron combines with a nuclear
proton to form a neutron and emit a neutrino.
5. Fission: A mode of radioactive decay in which a
nucleus of high mass number splits into two roughly equal and separate parts and, often, one or more
free neutrons.
6. Gamma Decay: A mode of radioactive decay in
which a high-energy photon (gamma ray) is emitted.
7. Ionizing Radiation: High-energy emission products of radioactive decay which ionize matter as
they pass through.
8. Isotope: Atoms having the same number of protons but different numbers of neutrons are isotopes
of one another. Deuterium is an isotope of hydrogen.
9. Mass Number: See Chapter 17.
10. Neutrino: An elementary (pointlike) particle emitted in beta decay. The neutrino is notable because
it lacks both a strong and an electromagnetic interaction with matter but interacts instead through the
weak interaction.
11. Neutron: A substructure of the nucleus of the
atom. Neutrons do not have an electrical charge. A
neutrons consist of three quarks. Protons and neutrons are both referred to as nucleons.
12. Proton: A substructure of the nucleus of the atom.
Protons are positively charged and consist of three
quarks.
13. Radioactive Dating: A method for measuring the
age of a sample by measuring the relative amounts
of radioactive elements and decay products in the
sample and accounting for the ratio in terms of the
number of half-lives that must have elapsed.
14. Radioactive Half-Life: A period of time in which
half the nuclei of a species of radioactive substance
would decay.
15. Radioactivity: Spontaneous changes in a nucleus
accompanied by the emission of energy from the
nucleus as a radiation.
Summary
The nucleus of each atom is a small, dense core
containing one or more protons and, with the exception
of 1H, one or more neutrons. The number of protons
(the atomic number) determines the chemical element
to which the atom belongs. Each element usually has
several isotopes—atoms with the same number of protons but different numbers of neutrons. Some of these
are unstable, or radioactive, and become more stable by
emitting ionizing radiation (an alpha particle, electron,
or electromagnetic radiation) or by fissioning. These
processes all release energy from atomic nuclei.
The rate at which radioactive nuclei decay is measured by their half-life and can be used to estimate how
much time has elapsed since certain kinds of events
took place. Much of our knowledge of the history of the
earth and its life forms comes from the study of radioactive materials and their by-products.
STUDY GUIDE
Chapter 24: The Nucleus
A.
1.
2.
3.
4.
5.
6.
trolled?
How can radioactive isotopes be used to determine
the date of an event?
FUNDAMENTAL PRINCIPLES
The Electromagnetic Interaction: See Chapter 4.
The Strong Interaction: See Chapter 2.
The Wave-Particle Duality of Matter and
Electromagnetic Radiation: See Chapters 14 and
16.
The Conservation of Mass-Energy: See Chapter
9.
The Conservation of Electric Charge: See
Chapter 7.
The Conservation of Mass Number: In radioactive decays, the number of nucleons (mass number) is conserved.
B. MODELS, IDEAS, QUESTIONS, OR APPLICATIONS
1. What are the parts of a nucleus and how is a nucleus described?
2. How does a nucleus spontaneously adjust to lower
energy arrangements?
3. Why must ionizing radiation be carefully con-
230
D. FOCUS QUESTIONS
1. Consider radioactivity:
a. Name and state the fundamental conservation
principle that accounts for the energy released in a
nuclear reaction. What is the source of the energy
released?
b. Write the equation describing beta decay in the
decay of carbon-14. Explain the meaning of the
equation.
c. What does half-life mean and what is the halflife of carbon-14?
d. Explain how carbon-14 is used to date events.
What assumptions are made and what limitations
are there in the analysis?
2. Consider radioactivity:
a. Name and state the fundamental conservation
principle that accounts for the energy released in a
nuclear reaction. What is the source of the energy
released?
b. Write the equation describing electron capture
in the decay of potassium-40. Explain the meaning
of the equation.
c. What does half-life mean and what is the halflife of potassium-40?
d. Explain how potassium-40 is used to date
events. What assumptions are made and what limitations are there in the analysis?
24.9. Iodine-131 is radioactive and emits beta rays.
What would be the atomic number and mass number of
the resulting nuclei? To which element would these
belong?
24.10. What is a positron? What happens to the
positrons that occur in nature?
24.11. What is radioactivity?
24.12. Mercury-195 is radioactive through the
process of electron capture. What would be the atomic
number, the chemical element, and the mass number of
the resulting nuclei?
24.13. Why would you expect fast-moving
charged particles from radioactive decay to cause ionization?
24.14. Why would you expect ions to behave differently from other atoms?
24.15. Suppose there are 100,000 radioactive
atoms in a sample of material. How many would be left
after one half-life has elapsed? Two half-lives? Three
half-lives?
24.16. Explain what is meant by ionization.
E. EXERCISES
24.1. A certain atom has a mass number of 40 and
an atomic number of 19. How many neutrons does it
have?
24.17. What is an ion?
24.18. Why would you expect radioactive decay to
be harmful to living systems?
24.2. How are the atoms of carbon-14 (radioactive
carbon) and carbon-12 (the usual form of stable carbon)
different from one another? How are they alike?
24.19. What is meant by the term “half-life”?
24.20. How can 14C be used to date an object?
24.3. Describe Rutherford’s evidence for the existence of the nucleus.
24.21. Explain how radioactive potassium can be
used to date a rock. What date or age is revealed by this
method?
24.4. Describe two kinds of particles of which
atomic nuclei are composed.
24.22. Review the description of carbon-14 dating.
What assumptions are being made that might affect the
precision of the method? What experiments could be
performed to reassure oneself that the assumptions are
valid?
24.5. What is meant by the atomic number of a particular nucleus?
24.6. What is meant by the mass number of a particular nucleus?
24.7. What is the difference between the various
isotopes of a given element?
24.23. List and describe three practical uses of ionizing radiation. What is a danger associated with ionizing radiation?
24.8. Radium-226 is radioactive and emits alpha
rays. What would be the atomic number and mass number of the resulting nuclei? To which element would
these belong?
24.24. Skeletal remains of a humanlike creature
were discovered in Olduvai Gorge in Tanzania in 1986.
The discoverers claim that the bones were found in a
geologic formation that is about 1.8 million years old.
231
Would carbon-14 dating be useful for establishing age
in this instance? Why?
24.25. Which of the following is true regarding the
isotopes of an element?
(a) They are equally radioactive.
(b) They have the same atomic mass.
(c) They have the same number of neutrons.
(d) They have the same number of protons.
(e) They have different numbers of protons.
24.26. In which of the decay processes is the atomic number of the final nucleus the same as that of the
original nucleus?
(a) alpha decay
(b) beta decay
(c) gamma deca
(d) electron capture
(e) fission
232