Fundamentals of Radiation & Radiation Safety
Fundamentals of Radiation & Radiation Safety Principles and Applications of Radiological Physics
Learning Objectives
At the end of this session, students should be:
Able to understand subatomic particles and explain atomic structure, including electron orbitals.
Capable of explaining the differences between elements, ions, isotopes, and compounds, including the types of molecular bonds.
Able to describe binding energy and what determines the binding energy of an electron and electronic configuration.
Converse on the structure of matter and its relevance to medical imaging.
Required Reading
Chapter 2: ‘Atoms and Matter’ from Graham's Principles and Applications of Radiological Physics, 7th Edition.
Introduction to Matter
Fundamental question: What is matter made of?
Atoms are identified as the fundamental building blocks of matter and have been the subject of extensive theoretical debate and experimental study.
Modern theories regarding atomic and subatomic structures are complex; however, many phenomena related to radiography can be explained with a relatively simple planetary model of the atom.
Historical Background of Atomic Theory
Etymology of the term 'atom': Originates from a Greek word meaning something which cannot be split.
Democritus (circa 400 BC): Proposed early ideas about particles.
Key Figures in Atomic Theory
John Dalton (1766 - 1844)
Developed the first atomic theory of matter following his experimental work on gases.
Proposed:
All elements are composed of tiny indestructible particles called atoms.
All atoms of a given element are identical and share the same atomic weight.
J.J. Thomson (1856 - 1940)
Discovered the electron during research on cathode rays.
Announced his findings in a lecture at the Royal Institute on April 30, 1897.
Eugen Goldstein (1850 - 1930)
Experimented with a perforated cathode ray tube, producing canal rays that travel in the opposite direction to cathode rays, concluding they must have an opposite charge.
Ernest Rutherford (1871 - 1937)
In 1911, concluded that atoms contain a tiny positively charged nucleus, with the majority of mass concentrated there and electrons occupying the space farthest from the nucleus.
Noted that atoms are mostly empty space but did not explain electron stability around the nucleus.
Niels Bohr (1884 - 1962)
Collaborated with Rutherford between 1911 and 1913, publishing a model where electrons orbit the nucleus in discrete energy levels, known as shells.
James Chadwick (1932)
Discovered a third type of subatomic particle, the neutron, with no charge, which helps reduce the repulsion between protons within the nucleus.
Atomic Structure and Models
Models of Atomic Structure:
Planetary Model: Describes a nucleus with orbiting electrons.
Quantum Model: Describes electrons as both particles and waves.
Subatomic Particles and their Properties
Particle | Symbol | Rest Mass (kg) | Rest Mass Energy (MeV) | Electric Charge |
|---|---|---|---|---|
Proton | p | $1.672 imes 10^{-27}$ | 938 | +1 |
Neutron | n | $1.675 imes 10^{-27}$ | 939 | 0 |
Electron | e | $9.109 imes 10^{-31}$ | 0.511 | -1 |
The atomic nucleus is very small and dense, containing protons and neutrons, typically ranging between $10^{-15}$ and $10^{-14}$ meters. The diameter of electron shells ranges between $1 imes 10^{-10}$ and $3 imes 10^{-10}$ meters.
Atomic Notation
Atomic Nomenclature: For an element represented as EA Z:
E = element name
A = atomic mass number (total of protons and neutrons)
Z = atomic number (number of protons)
Nuclides and Isotopes
Nuclide: An atom characterized by a specific number of neutrons.
Example isotopes of carbon:
C^{12}_{6} (6 protons + 6 neutrons)
C^{13}_{6} (6 protons + 7 neutrons)
C^{14}_{6} (6 protons + 8 neutrons)
All nuclides of an element share the same atomic number (number of protons).
Stability of the Nucleus
Atomic nuclei consist of positively charged protons and neutral neutrons.
Nuclei should theoretically explode due to electrostatic forces, but many nuclei are stable due to the balance of electrostatic forces and short-range nuclear forces.
The energy maintaining nuclear stability is referred to as nuclear binding energy.
Electron Orbitals
Electron Shells:
Electrons reside in orbital shells around the nucleus.
In stable atoms, the number of electrons equals the number of protons, resulting in electrical neutrality.
Energetics of Electron Configuration:
Electrons fill the inner shells first (lowest energy state).
Quantum Perspective of Electrons
Wave Function: Mathematical description providing the probability density of electron locations across energy levels.
Predicting precise electron locations is not possible; only probabilities of locations are assessed based on mathematical models.
Electron Energy Levels
Formula for maximum number of electrons in a shell:
ext{Max electrons per shell} = 2n^2 where $n$ = shell number.
Details of shell configuration:
Shells:
K (n=1): 2 electrons
L (n=2): 8 electrons
M (n=3): 18 electrons
N (n=4): 32 electrons
Binding Energy
Electron Binding Energy: The work needed to remove an electron from an atom’s energy level.
Expressed in electron volts (eV) or kilo-electron volts (keV); $1 ext{ keV} = 1000 ext{ eV}$, and $1 ext{ eV} = 1.6022 imes 10^{-19} ext{ joules}$.
Factors Affecting Binding Energy:
Atomic Number: Higher atomic number results in higher binding energy due to increased proton count in the nucleus.
Nucleus-Electron Distance: Greater distance results in decreased binding energy due to the reduced positive pull from the nucleus.
Pauli Exclusion Principle: No two electrons can occupy the same quantum state simultaneously.
Ionization and Excitation
Excitation
Occurs when an electron absorbs energy sufficient to promote it to a higher energy level without overcoming binding energy.
Atoms in an excited state release energy and return to ground state, emitting electromagnetic radiation.
Ionization
Ionization occurs when an electron absorbs enough energy to overcome its binding energy, escaping the atom and resulting in a charge imbalance (net positive charge).
Ions are defined as atoms with unequal numbers of protons and electrons.
Molecular Compounds and Bonding
Types of Compounds
Molecular Compounds: Made from atoms of more than one element chemically bound in fixed ratios.
Examples: ( O2 ) (Oxygen) and ( CH4 ) (Methane).
Types of Bonds
Ionic Bonding:
Atoms with nearly full or empty outer electron levels may become chemically reactive.
Example: Sodium (atomic number 11) donates an electron to Chlorine (atomic number 17), forming NaCl through ionic bonding.
Covalent Bonding:
Atoms with half-full outer shells share electrons to achieve stability.
Examples: ( O2 ) (Oxygen) and ( CH4 ) (Methane).
Other Types of Matter
Crystalline Compounds: Atoms in a fixed lattice; example includes caesium iodide (CsI) used in x-ray detectors.
Inert Gases: Monoatomic and chemically unreactive due to completely filled outer electron shells.
Example: Helium is utilized in superconductive MRI scanners.
Metals: Exhibit luster, malleability, and conductivity due to metallic bonds with delocalized electrons.
States of Matter
Common States: Gases, Liquids, Solids.
State transition is influenced by temperature:
Gases: No intermolecular bonds.
Liquids: Weak intermolecular bonds.
Solids: Strong intermolecular bonds.
Radiographic Representation:
Low-density materials appear black; high-density materials appear white on x-ray images.
Bodily tissues exhibit shades of grey based on density and thickness.
Summary
An atom consists of neutrons (0 charge) and protons (+1 charge) forming the nucleus, collectively referred to as nucleons.
Mass characteristics:
Neutron: mass of 1, charge of 0.
Proton: mass of 1, charge of +1.
Electron: mass of 0.0005, charge of -1.
Mass Number (A): Total of protons and neutrons.
Atomic Number (Z): Total number of protons.
Maximum number of electrons per shell is calculated by the formula: 2n^2 .
Binding energy of electrons corresponds to their position relative to the nucleus and overall atomic number.