Nuclear Medicine Physics
Course No. 311RAD
Dr. Mohammed Alnafea
alnafea@ksu.edu.sa
www.ksursd.net
Course Objective:
Physics and instrumentation impact all of
the sub-specialty area of nuclear
medicine because of their fundamental
importance.
This course is designed to provide an
introductory course to the two
subsequent nuclear medicine courses
namely (RAD 432 and RAD 466).
Aim of the Course:
This course aims to give the student the
following:
1. A review of the general physical
principle of nuclear medicine, its
procedures and instrumentations to be
of permanent value to them to be able
to perform the basic nuclear medicine
procedures.
Aim of the Course:
2. A broad overview of the basic
elements contained in the course
as well as gives students an
orientation to the operation of a
nuclear medicine department.
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Course Content
• The course will cover several topics
including atomic and nuclear structure,
production of radioisotopes,
radioactive decay, the radioactive
decay law, unit of radiation
measurements, attenuation of
gamma-rays, scintillation detectors,
nuclear medicine imaging systems and
radiation protection in nuclear
medicine.
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Course Contents
– Topics:
» Atomic & Nuclear Structure
» Radioactive Decay
» The Radioactive Decay Law
» Production of Radioisotopes
» Units of Radiation Measurement
» Interaction of Radiation with Matter
» Radiation Protection in Nuclear
Medicine.
» Attenuation of Gamma-Rays
» Scintillation Detectors
» Nuclear Medicine Imaging Systems6
Course References:
1. Medical imaging physics, William R. Hendee, E. Russell
Ritenour, 4th edition, 2002.
2. Physics in Nuclear Medicine, James A. Sorenson (Editor),
Michael E. Phelps (Editor) Grune & Stratton; ISBN:
0808918044; 2nd edition (January 1987).
3. An Introduction to the Physics of Nuclear Medicine by Paul N.
Goodwin Charles C Thomas Pub Ltd; ISBN: 0398035695; (June
1977).
4. Essentials of Nuclear Medicine Physics
by Rachel A., Md Powsner, Edward R. Powsner Blackwell
Science Inc; ISBN: 0632043148; 1st edition (September 15,
1998).
5. Essentials of Nuclear Medicine Imaging by Fred A., Jr., M.D.
Mettler, Milton J., Md. Guiberteau, Barbara Mettler W B
Saunders Co; ISBN: 0721651216; 4th edition (January 15,
1998)
Evaluation & Assessment Methods
Students are assessed as follows:
1.
1st mid term exam
10 marks
2.
2nd mid term exam
10 marks
3.
Assignment or report
10 marks
4.
Oral Presentation
10 marks
5.
20 marks
6.
Hospital & Practical
report + Oral exam
Final exam
7.
Total
100 marks
40 marks
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Academic year: 1429-1430H
Week 1
Week 2
Week 3
Week 4
Week 5
1st mid term exam
Week 6
Week 7
Week 8
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Academic year: 1429-1430H
Week 9
Week 10
Assignment or report
Week 11
2nd mid term exam
Week 12
Oral Presentations
Week 13
Hospital report + Oral exam
Week 14
Week 15
Week 16
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Oral Presentation
– A list of topics of the oral presentations for the
course RAD-311
1. The Physical and chemical properties of NaI
detector.
2. Detectors & new detector materials.
3. Radionuclide for PET and their production.
4. Baby cyclotron for PET.
5. 2 & 3  annihilation and their uses in PET.
6. The principle of planar gamma camera
7. The principle of SPECT imaging.
8. SPECT in Brain Function Imaging.
9. Molecular imaging with PET & SPECT.
10.Small animal imaging.
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11.FDG and blood flow.
Atomic & Nuclear Structure
The Atom is the smallest indivisible
particle of matter, i.e. something
that cannot be divide.

All atoms are composed of just
three particles: protons, neutrons &
electrons.
The centre of an atom is called the
nucleus.
Atomic & Nuclear Structure
The nucleus contains protons (P) & neutrons (N).
P & N have virtually the same mass i.e. same as the
hydrogen atom. So P & N are given a relative mass of
one like the hydrogen atom.
Protons are Positively charged, but neutron are
neutral.
Atomic & Nuclear Structure
» More than 99.9% of the atom is empty space
occupied by moving electrons (E).
» E have a mass about 2000 times less than P &
N.
» Electrons are negatively charged.
» Negative charge on one electron just balances
the positive charge on one proton.
» The electrons move around the nucleus very
rabidly and remains in layers or shells at
different distances from the nucleus.
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Atomic Structure
Particle
Position
Mass
with the relative to a
atom
hydrogen
atom
Proton
nucleus
Neutron nucleus
Electron
Shells
Charge
Mass (gm)
1
+1
1.672810-24
1
0
1.67410-24
0.0005
-1
9.0910-27
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Atomic Structure
» Protons, neutrons & electrons are the building blocks for
atoms.
» Different atoms have different number of protons,
neutrons and electrons.
» Hydrogen atoms are the simplest atoms as each atom
contains one proton & one electron.
» Some of the heaviest atoms have 100 or more protons,
neutrons & electrons. For instance, each atom of lead
has 82 protons, 82 electrons and 125 neutrons.
» NB: Why same No. of protons and electrons? This allows
the +ve charges in the protons to balance the –ve
charges on the electrons so each atom is neutral overall.
» The electrons move around the nucleus very rabidly and
remains in layers or shells at different distances from
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the nucleus.
Atomic number & mass
number
12
6
C

The No. of protons in an atom tells you which elements it is and
thus, is called atomic number (Z).

Thus, hydrogen has atomic number of one.

Notice that the mass of an atom depends on the No. of protons &
neutrons in its nucleus.


Atomic number (Z) = No. of protons (P)
Mass number (A) = No. of protons (P) + No. of neutrons (N).

The mass No. (A) & atomic No. can be shown with the symbol for
an element.

e.g. carbon atoms with 6 protons & 6 neutrons can be represented
as
mass No.( A) 12
atomic No. ( Z ) 6
C
17
12
6
C
Relative atomic mass & isotopes

Several elements have
relative atomic masses
which are whole numbers.
For example Al have:
13 P- relative mass =13
13 E- relative mass =0.0
14 N- relative mass =14
--------------------------total relative mass =27

Relative mass of an atom
sometimes referred as
one mole of the element.
Element
Symbol
relative
mass
Carbon
C
12.0
Hydrogen
H
1.0
Oxygen
O
16.0
Copper
Cu
63.5
18
12
6
C
Isotopes

Some elements have relative atomic masses that
are not whole numbers e.g. Iron (Fe) =55.8

One element could have atoms with different
masses. Atoms of the same element with
different masses are called isotopes. For e.g.
naturally occurring chlorine Cl has two isotopes
35
17
cl
37
17
cl

Each of the above isotope has 17 protons & 17
electrons. Both have atomic No. of 17. but one
has 18 neutrons & the other has 20 neutrons.

Isotopes are atoms with the same atomic
No., but different mass numbers.
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Isotopes of Hydrogen
• The most common
isotope of hydrogen
has no neutrons at all;
• there's also a
hydrogen isotope
called deuterium,
with one neutron, and
another, tritium, with
two neutrons.
Hydrogen-1
Hydrogen
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium)
1H
2H
3H
Z=1
A=1
Z=1
A=2
Z=1
A=3
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Other forms of atoms
• Isobars are nuclides having the same mass
number, i.e. sum of protons & neutrons;
Carbon-12 & Boron-12.
• Isotones two nuclides are isotones if they
have the same number of neutrons N. For
example, Boron-12 and Carbon-13 both have
7 Neutrons.
• Nuclear isomers are different excited states of
the same type of nucleus e.g. 99mTc (has
different amount of nuclear energy.
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Electron Structures
& the Periodic Table
1st lecture Questions
1.
Show how the mass No. & atomic No. can be represented for
sodium and aluminium?.
Na has 11 P & 12 N but Al has 13 P & 14 N.
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11
2.
Na
27
13
Al
Calculate the relative atomic mass of naturally occurring chlorine
contain three chlorine-35 atoms for every chlorine-37.
RAM= (335)+37/4=35.5
3.
There are three isotopes of the element hydrogen what
are they and how do we distinguish between them?
Hydrogen has two naturally occurring isotopes, normal
hydrogen, which contains no neutrons, and heavy hydrogen,
which has one neutron in each atom. The third isotope,
tritium, has two neutrons.
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Electron Structures & Periodic
Table
During chemical reactions, changes occur in the No. of
electrons belonging to atoms.
 Some atoms gain E, some lose E & other share E. NB:
once these changes in E structure have taken place, the
atoms or ions are usually more stable i.e. less reactive.
Some atoms never react e.g. helium & neon. Why? Have
very stable electron structures.
Atom or ions will have very stable electrons structure if
they have 2 E (like helium), 10 E (like neon) 18 E (like
argon) & so on. These stable E structures are closely
related to the way in which E are arranged in layers or
shell.
The No. of E in an atom in each shell is determined by this
equation:
No. of electrons= 2 n2
Electron Structures & Periodic
Table
The No. of E in an atom in each shell is determined by this equation:
No. of electrons= 2 n2
Where n is the shell number i.e. 1 for K, 2 for L, 3 for M…..
For example Na has 11 protons:
=2
=8
at (K shell) n=1 the No. of E = 2(1)2
at (L shell) n=2 the No. of E = 2(2)2
at (M shell) only one E
The above equation tells us that the capacity of the 1st shell is 2
electrons and the capacity of the 2nd shell is 8 electrons. The 3rd
shell has a total capacity of 2 x 32 = 18 electrons
 NB: E shells actually have sublevels. The first sublevel (the s
sublevel) holds 2 electrons. The second, p, sublevel holds 6. The
third, d, sublevel holds 10. After levels 3s and 3p are filled, E shell
#3 acts as if it has reached capacity with only 8 total E.
 In other words, in an atom with 20 E (which is the element
calcium, Ca) the first 2 electrons are located in the 1st shell, the next
8 in shell #2, the following 8 in shell #3 and the remaining 2
2Al
1
7
3
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Electron Structures & Periodic
Table
The atoms can be arranged according to their electron structure
and chemical properties. This is where we encounter the Periodic
Table of Elements.
The periodic "law" of chemistry recognises that many properties of
the chemical elements are periodic functions of their atomic number
(the number of protons within the element's atomic nucleus).
 The periodic table is an arrangement of the chemical elements
ordered by atomic number in columns (groups) and rows (periods)
presented so as to emphasize their periodic properties.
 Atoms are ordered by their atomic number in the Periodic
Table. The Table is set up so as to indicate the number of electron
shells in each atom and the number of valence electrons (electrons
in the outermost shell) in the atom.
 As you read the Table from left to right in any one row, the
number of valence electrons increases. For example, hydrogen has
1 electron (in the first shell). Helium (He), the 2nd element in the
first row, has 2 electrons (thus filling its valence shell).
Electron Configuration Table

The electron configuration is the arrangement of electrons in an
atom, molecule, or other physical structure. It basically describes how many
electrons are in each energy level of an atom and how the electrons are
arranged within each energy level.
For example, this is the electron configuration table for gold (Au):
Binding Energy
• In general, binding energy represents the mechanical work
which must be done in acting against the forces which hold an
object together, while disassembling the object into
component parts separated by sufficient distance.
• Nuclear binding energy is derived from the strong nuclear
force and is the energy required to disassemble a nucleus into
free unbound neutrons and protons, strictly so that the
relative distances of the particles from each other are infinite
(far enough).
• Nuclei are made up of protons and neutron, but the mass of a
nucleus is always less than the sum of the individual masses
of the protons and neutrons which constitute it.
• The difference is a measure of the nuclear binding energy
which holds the nucleus together.
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Binding Energy
• If a composite object is stable, for instance, the nucleus of
a helium atom does not spontaneously split into the two
protons and two neutrons that are its constituents:
• Energy of the composite object + energy expended to split
it up = sum of the energies of the separate parts after the
split.
• We can also move the expended energy to the right side of
the equation, which leaves us with
•
Energy of the composite object = sum of the energies of
its parts - energy needed to split the object apart.
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Binding Energy
• Nuclear binding energy can be calculated from the Einstein
relationship:
Nuclear binding energy = Δmc2
• According to Einstein, to every energy there corresponds a
mass, and to every mass there can be assigned a
corresponding energy.
• Einstein showed that mass and energy are really two
different forms of the same thing; the "vanishing" mass of
the protons and neutrons is simply converted to energy.
• For most nuclei, the binding energy per nucleon is about 8
MeV.
• The relativistic mass of a bound system is somewhat
smaller than the sum of the masses of its constituent parts,
namely:
• Mass of bound system = sum of masses of its parts 22/03/2016
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2.lecture
(binding energy)/c
Binding Energy
• The unit for binding energies per
nucleon is the Mega-electron-Volt
(MeV).
• One MeV is defined as the energy
gained by an electron accelerated
by an electrical voltage of one
million Volt (it is also the energy
corresponding to twice the mass of
an electron at rest).
• The comparison of the alpha
particle binding energy with the
binding energy of the electron in a
hydrogen atom is shown in Fig.
• The nuclear binding energies are on
the order of a million times greater
than the electron binding energies
of atoms.
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Electron Binding Energy
• Electron binding energy is a measure of the energy required to
free electrons from their atomic orbits.
• The unit is electron volt (ev).
• Electron volt- the energy required by an electron to be accelerated
through 1 volt potential difference
• Electrons can only have discrete energy levels i.e. K, L, M, …. shells.
• Electrons in the K-shell are tightly bound electrons
• K shell has maximum energy (due to its stability).
• To remove an electron from its shell E  electron binding
energy.
• NB: Binding energy increases rapidly with atomic No. (Z).
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Ionization
• Ionization is the physical process of converting an atom or
molecule into an ion by adding or removing charged particles
such as electrons or other ions.
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Ionization
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Excitation
• Excitation is an elevation in energy level above an
arbitrary baseline energy state.
• Excitation by absorption of light and de-excitation by
emission of light
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Questions
1.
Given that 1u has a mass of 1.65511×1027kg, calculate the
energy (MeV) of 1u using Einstein's equation?.
2.
From the graph
calculate the binding
energy per nucleon for
the following
elements:
Iron (Nucleon number
- 56)
a.
b.
Magnesium (Nucleon
number - 24)
c.
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Titanium
(Nucleon
number - 48)?.
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RAD 311
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Answer
1.
Given that 1u has a mass of 1.65511×10-27kg, we can calculate
the energy of 1u.
E = mc2
= (1.65511×10-27) × (3.0×108)2 J
= 1.489599 × 10-10J
However, 1eV = 1.6×10-19J, so
E= 1.48959 × 10-10 / 1.6×10-19 eV= 9.31 × 108 eV= 931 MeV
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Questions
1.
Show how the mass No. & atomic No. can be
represented for sodium and aluminium?.
2.
Calculate the relative atomic mass of
naturally occurring chlorine contain three
chlorine-35 atoms for every chlorine-37.
3.
There are three isotopes of the element
hydrogen what are they and how do we
distinguish between them?
40