Chapter
NP-1
Nuclear Physics
Atomic Nature of Matter
TABLE OF CONTENTS
INTRODUCTION
OBJECTIVES
1.0
PROPERTIES OF SUBSTANCES
1.1
2.0
CHEMICAL AND PHYSICAL PROPERTIES
COMPOSITION OF ATOMS
2.1
ATOMIC STRUCTURE
2.2
NUCLEAR NOTATION
2.3
ATOMIC MASS AND WEIGHT
2.4
ELECTRON STRUCTURE
3.0
THE NUCLEUS
4.0
ISOTOPES
4.1
5.0
ISOTOPIC ABUNDANCE
SUMMARY
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INTRODUCTION
This is the first of five lessons in nuclear physics. This lesson, will present the basic
atomic nature of matter. Included in the discussion is information on Atomic Structure,
Atomic Mass and Atomic Weight. At the completion of this lesson, the Contractor Health
Physics Technician should have a good understanding of the construction of an atom, and
the standard nomenclature for describing atoms. The information presented in this lesson
provides the foundation for ideas that are more complex.
OBJECTIVES
TERMINAL OBJECTIVE
The Contractor Health Physics Technician will describe the structure and properties of
atoms and their constituents. The Contractor Health Physics Technicians will describe the
idea of ionization.
ENABLING OBJECTIVES
Upon completion of this lesson, the Contractor Health Physics Technician will be able to:
1. Describe the components of atoms including each component's symbol, relative mass,
and relative electrical charge.
2. Recognize and manipulate atomic symbols to determine the number of protons,
neutrons, or chemical symbol.
3. Define isotope.
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1.0
O F
M A T T E R
PROPERTIES OF SUBSTANCES
Atoms are the basic building blocks of all matter. There are presently 105 elements (92
occur naturally, 13 made artificially), each element having atoms with the same number
of protons, but possibly with different number of neutrons. The atom is the smallest
particle of an element that can enter into a chemical reaction. Figure NP-1-1 lists 22
elements common to nuclear power along with their symbols.
1.0
ELEMENTS COMMON TO NUCLEAR PLANTS
ELEMENTS
SYMBOL
ELEMENTS
SYMBOL
ANTIMONY
Sb
LEAD
Pb
BERYLLIUM
Be
LITHIUM
Li
BORON
B
NICKEL
Ni
BROMINE
Br
NITROGEN
N
CARBON
C
OXYGEN
O
CHLORINE
Cl
PLUTONIUM
Pu
CHROMIUM
Cr
SAMARIUM
Sm
COPPER
Cu
SODIUM
Na
HELIUM
He
URANIUM
U
HYDROGEN
H
XENON
Xe
IRON
Fe
ZIRCONIUM
Zr
Figure NP-1-1 ELEMENTS COMMON TO NUCLEAR PLANTS
When atoms combine and join, they form molecules. A molecule is the smallest particle
of a substance that still retains the characteristics of that substance. In many cases, a
single atom displays the properties of the substance, so that a molecule consists of a
single atom. This is true for metals such as iron, lead, aluminum, etc. Molecules of
compounds consist of one or more atoms of at least two elements in combination. The
properties of a substance depend on the structure of the molecule and frequently do not
resemble the properties of the elements from which they are produced. Sodium (an
explosive metal) and chlorine (a poisonous gas) combine to form sodium chloride,
common table salt.
1.1
CHEMICAL AND PHYSICAL PROPERTIES
Chemical properties are those which deal with a substance entering into a reaction with
other substances. These properties include the ability to burn or corrode and form
compounds.
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Chemical reactions concern the transfer or sharing of outer shell electrons, as shown in
Figure NP-1-2 and NP-1-3.
Chemical reactions do not involve the atom's nucleus and the energy involved in
chemical reactions is about a million times less than nuclear reactions.
Physical properties are those which can be observed without a chemical change taking
place in the material. Some physical properties are: color, density, hardness, elasticity,
melting point, boiling point and crystalline structure.
NP-1-2 Formation of Sodium Chloride
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NP-1-3
Formation of a Water Molecule
2.O COMPOSITION OF ATOMS
Atoms, the building blocks of all
substances, contain even smaller
(sub-atomic) particles. All atoms
consist of a positively charged
nucleus surrounded by a number
of negatively charged electrons.
An adequate model for our
purposes includes electrons as
very small particles revolving
around a central nucleus in
orbits (similar to planets
revolving around the sun).
Figure NP-1-4 illustrates the
model.
NP-1-4
Composition of Atoms
Atoms are normally electrically neutral because the number of electrons equals the
number of protons. The coulombic force (opposite charges attract) attracts electrons
toward the nucleus. They continue to orbit the nucleus because of the kinetic energy they
possess. They can escape from the atom only if they obtain sufficient energy to overcome
the forces that hold them in orbit.
Atoms are mostly empty space; the nucleus has a diameter of about 10-14 meters and the
entire atomic radius is about 10-10 meters. The nucleus contains about 99.97% of the mass
of the atom in about 10-10% of the volume. If you could increase the size of the nucleus to
that of a baseball, the distance between it and the nearest orbital electron would be about
one-half mile!
2.1
ATOMIC STRUCTURE
The three basic particles that make up an atom are the proton, neutron and electron. A
description of each particle's physical properties is as follows:
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PROTON
The proton is an elementary particle located in the atom's nucleus. It has a
positive electrostatic charge and a mass of 1.6724 x 10-24 grams, symbol p.
NEUTRON
The neutron is an elementary particle also located in the atom's nucleus. It has no
charge. Its' mass is slightly greater than the proton at 1.6747 x 10-24 grams.
Presently the character η is in use to symbolize it.
ELECTRON
The electron is an elementary particle that orbits the nucleus. It has a negative
electrostatic charge, equal in charge intensity to the proton. Its mass is 1/1838 that
of a proton, symbol e.
An atom consists of the protons and neutrons tightly bound in a central nucleus,
surrounded by the electrons orbiting in "shells" or energy levels.
The nucleus carries a positive electric charge due to the presence of protons, so that the
total charge is the sum of the total number of protons. This positive charge balances the
total negative charge of the orbiting electrons, so that the whole atom is electrically
neutral. Thus, the number of orbital electrons equals the number of protons in the nucleus
for a neutral atom. The number of electrons determines the atom's chemical properties,
which are largely uninfluenced by the nucleus itself.
2.2
NUCLEAR NOTATION
Figure NP-1-5 illustrates a simple method for designating any nuclide. This accepted
notation is in common usage.
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Recall again the definitions of atomic mass number and atomic number.
Ø Atomic number (Z) is the number of protons in the nucleus. It defines the
element.
Ø Atomic mass number (A) is the sum of the number of protons and neutrons in the
nucleus.
Therefore, the number of neutrons in a given nucleus is the difference between the atomic
mass number and atomic number.
Number of neutrons = A – Z
Hydrogen, helium and uranium atoms are examples of this designation and are as
follows:
1H
1
4
2He
92U
235
0 neutron
=
1 proton
=
1 proton H
=
2 protons
+
2 neutrons
=
92 protons
+
143 neutrons =
=
2 neutrons
2 protons He
92 protons U
143 neutrons
To determine the number of particles in any atom, the following rules apply:
Ø The number of protons in the nucleus is always equal to the atomic number of the
element (Z).
Ø The number of neutrons is equal to the difference between the atoms atomic mass
number and atomic number (A-Z).
Ø The number of orbital electrons is equal to the number of protons in electrically
neutral atoms.
2.3
ATOMIC MASS AND WEIGHT
Since atomic masses are about 10-24 grams, a need for a more convenient unit exists.
The chosen standard is the "atomic mass unit" (AMU). One AMU is equal to 1/12 the
mass of a carbon 12 nucleus (carbon 12 contains six protons and six neutrons). Present
evidence indicates that 1 AMU is very nearly the mass of a proton or neutron. The
accurately determine value is:
1 AMU = 1.6605 x 10-24 gm
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The atomic weight of an atom is its mass in AMUs. The nuclear mass differs from the
atomic weight by the total masses of the electrons. Further, the atomic weight of an
element must be considered in detail. A naturally occurring element consists of all those
atoms whose nuclei contain a given number of protons, which is the atomic number. For
example, the element carbon contains all naturally occurring atoms with atomic number
6. However, the atomic number does not completely describe the nucleus, since there
may be differing numbers of neutrons included in all nuclei with an atomic number of 6.
Each Nuclei of a given element, has the same atomic number, but possess different
numbers of neutrons, and are known as the isotopes of the element.
The electron structure of an atom determines the chemical properties of an atom. The
isotopes of an element make up a set of different atoms having the same chemical nature,
but differing in atomic weight. Thus, the atomic weight of an element is the average
atomic weight of all the naturally occurring isotopes of that element. Taking into account
the relative isotopic abundance, most elements consist of 2 or more isotopic types.
Isotopes will be discussed in greater detail in Section 4.0
2.4
ELECTRON STRUCTURE
The configuration of the orbiting electrons determines the chemical properties of an
atom. Specifically, these properties depend upon the number of electrons in the outermost
shell (valence electrons). Each shell represents a different energy level or energy state,
the innermost shell representing the lowest energy state -- the outermost shell the highest.
If a shell has less than the allowed number of electrons, an electron from a higher energy
shell can "drop" into the vacant spot (with X-rays or visible light range energy given off
in the process). An atom is in the "ground" state when its electrons are in the innermost or
lowest energy level shells. An atom in an "excited" state has an electron in a higher
energy shell.
Each shell has a maximum number of electrons that it will hold. Table 1 gives the
maximum number of electrons for any given shell.
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Shell Designation
K
L
M
N
O
P
Q
O F
M A T T E R
Table 1 – Maximum Number of Electrons per Shell
Number (n) of shell Maximum Number
from center
of Electrons in Shell
1
2
2
8
3
18
4
32
5
32
6
12
7
2
Total Electrons
2
10
28
60
92
102
106
Table 1 demonstrates that if there are 2 electrons they will occupy the K shell. In an atom
with more than 2 and less than 10 electrons, the first 2 electrons are in the K shell and the
remaining electrons are in the L shell. In atoms with more than 10 but less than 28
electrons, the first 2 electrons are in the K shell, the next 8 are in the L shell, and the
remaining electrons are in the M shell, etc... . However, the outermost shell (valence
shell) cannot contain more than 8 electrons. Atoms with only 1 or 2 valence shell
electrons are highly reactive in chemical reactions. Atoms with a full valence shell
normally do not react at all. See Figures NP-1-6 and NP-1-7.
NP-1-6 Electron Structure
NP-1-7 Electron Structure
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3.0
O F
M A T T E R
THE NUCLEUS
As mentioned earlier, the atomic nucleus contains protons and neutrons closely packed
together. Clearly some force must be present that holds the nucleus together and
overcomes the strong coulombic repulsive forces of the tightly packed protons. The
"nuclear force" is the name for this very strong force. The nuclear force acts between
fundamental particles that are neither electrical nor gravitational in nature. The nuclear
force is a "short-range" force, and when protons or neutrons are within about 10-13cm
the nuclear force binds them together strongly, overcoming any electrostatic repulsion
between protons.
Nuclear forces have several important characteristics:
Ø They are charge independent; that is, the force is just as strong whether neutrons
or protons are being considered.
Ø They are extremely strong, much stronger than gravitational or electrical forces.
Ø They have very short range, about 10-13 cm.
Ø They are saturable; one particle can only exert a nuclear force on a limited
number of other particles.
Although nuclear forces are much stronger than electrostatic forces, the electrostatic
forces play a role in tending to force the nucleus apart. Because neutrons can add nuclear
force without electrostatic repulsion, their presence in the nucleus is very important. This
gives rise to the following two important facts:
1.) there is no stable nuclei consisting of two or more protons with no neutrons
2) as the number of protons in a nucleus becomes greater, a relatively greater
number of neutrons needed to stabilize the nucleus.
The "belt of stability" shown in Figure NP-1-8 (a plot of stable nuclei) demonstrates the
latter point. (The terminology used here is that an "unstable" nucleus is a combination of
protons and neutrons that will not remain a unit indefinitely. An unstable nucleus emits
energy or particles in a process called "radioactive decay".)
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Each box in Figure NP-1-8 represents a known nucleus containing Z number of protons
and n number of neutrons. The straight line represents the line along which nuclei would
lie if they contained equal numbers of protons and neutrons. For nuclei with mass
numbers less than 40, the stable isotopes contain approximately equal numbers of protons
and neutrons. However, the heavier stable isotopes contain considerably more neutrons
than protons. The "belt of stability" would encompass a line drawn through the average
location of stable nuclei on this graph. Observations indicate that nuclei on either side of
the line have too low or too high a N/P ratio for stability and will undergo radioactive
decay to bring the nucleus closer to stability.
NP-1-8 Nuclei for Z and n
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4.0
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ISOTOPES
Not all atoms of a particular element are exactly alike. Hydrogen, for example, exists in
three forms. The most abundant form of hydrogen exists as one proton and one electron.
The deuterium atom, an isotope of hydrogen, differs from ordinary hydrogen in that
deuterium contains a neutron together with the proton in the nucleus. Tritium, another
isotope of hydrogen, has two neutrons and one proton in the nucleus. It follows that the
three isotopes of hydrogen differ in their atomic mass. Figure NP-1-9 illustrates the three
isotopes of hydrogen.
We can the define isotopes as: atoms that have the same atomic number, but different
atomic weights.
To identify individual isotopes of an element, use the appropriate mass number; e.g.,
hydrogen-1, hydrogen-2 and hydrogen-3
identify the isotopes of hydrogen.
Most elements exist in nature as a
mixture of two or more isotopes. For
example, there are two stable isotopes of
NP-1-9
Isotopes of Hydrogen
nitrogen with mass numbers 14 and 15,
as shown in Figure NP-1-10.
NP-1-10
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These isotopes of nitrogen differ only in the fact that N-15 has one more neutron in the
nucleus. The electron structure is the same, otherwise the two isotopes would have
different chemical properties.
Since all isotopes of an element have the same number of protons in the nucleus and the
same number and distribution of electrons, we can properly define an element as a
substance consisting entirely of atoms having the same atomic number; that is, the same
number of protons.
4.1
ISOTOPIC ABUNDANCE
Isotopic abundance refers to the percentage of each isotope in an element's natural state.
For instance, ordinary copper has an abundance of 69.2% of copper-63 and 30.8% of
copper-65. The composition of natural silver is of 51.83% of silver-107 and 48.17%
silver-109. Carbon has an isotopic abundance of 98.89% carbon-12 and 1.11% carbon13.
Definition:
Isotopic abundance is the amount of the isotope (percentage)
present in a normal natural mixture of the element.
As an example of an element existing in a natural state, consider uranium. Natural
uranium consists of a mixture of isotopes having mass numbers 234, 235 and 238. Table
2 lists the isotopic abundance of natural uranium isotopes.
Table 2
Composition of Natural Uranium
Isotope
Percent Abundance
U-234
0.006%
U-235
0.714%
U-238
99.28%
Natural U
Atomic Mass
234.0409
235.0439
238.0508
238.03
The atomic mass given in Table 2 for natural uranium is a weighted average based on the
percent abundance of the atomic mass of each isotope. It is the atomic mass given for
uranium on any modern periodic chart. This is the case for all atomic masses given in
periodic tables in that the table gives the atomic mass of the natural element.
5.0
SUMMARY
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Atoms consist of 3 basic particles: protons, neutrons and electrons. The protons and
neutrons located in the nucleus about which the electron(s) orbit. The nucleus contains
almost all the mass of the atom, but occupies only a very small part of the atom's entire
volume. Atomic mass, atomic weight, and the atomic mass unit (AMU), were the
subjects of this lesson. The AMU is equal to 1/12 the mass of a carbon-12 nucleus.
All the atoms of a particular element have the same number of protons (Z). However, not
all atoms of an element have the same atomic mass (A). Atoms with the same Z, but
different A numbers are isotopes. The following notation can identify an isotope: ZXA,
element name, or symbol and atomic mass number. For example, 92U235 or uranium-235
or U-235.
Electrons orbit the nucleus in "shells". Each shell represents a discrete energy level. Each
shell has a limit on the number of electrons it may contain. The shell closest to the
nucleus (K shell) represents the lowest energy state, and the outermost shell, the highest
energy state.
The nucleus is a tightly bound cluster of neutrons and protons called nucleons. The
"nuclear force" holds the nucleons within the nucleus. This force is charge independent,
extremely strong, has a very short range, and is saturable. Although nuclear forces are
much stronger than electrostatic forces, the electrostatic forces tend to force the nucleus
apart. Neutrons add nuclear force without adding electrostatic repulsion.
Current theory is that the stability of the nucleus is dependent on the neutron to proton
ratio of the nucleus. A nucleus with too low or too high of an n/p ratio for stability will
undergo radioactive decay.
REFERENCES
Freeman, Ira M. Physics Made Simple, Garden City, N.Y.: Doubleday and Co., 1965.
Shortly, George; William Dudley. Principles of College Physics, Second Edition.
Englewood Cliffs, N.J.: Prentice-Hall, Inc., 1967.
Tulley, Donald E.; Thumm, Walter. Physics for College Students with Applications to the
Life Sciences, Menlo Park, CA: Cummings Publishing Co., 1974.
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LIST OF FIGURES
NP-1-1
NP-1-2
NP-1-3
NP-1-4
NP-1-5
NP-1-6
NP-1-7
NP-1-8
NP-1-9
NP-1-10
Chemical Symbols
Formation of Sodium Chloride
Formation of a Water Molecule
Composition of Atoms
Nuclear Notation
Electron Structure
Electron Structure
Nuclei for Z and n
Isotopes of Hydrogen
Isotopes of Nitrogen
LIST OF TABLES
Table 1 - Maximum Number of Electrons per Shell
Table 2 - Composition of Natural Uranium
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