1.2 Fundamental Concepts and Terminology

Atoms, molecules, and subatomic particles form the foundation of atomic physics. Understanding these building blocks is crucial for grasping how matter behaves at the smallest scales.

The periodic table organizes elements based on their properties, while isotopes and ions add complexity to atomic structure. Mastering atomic notation helps us describe and analyze these fundamental components of matter.

Essential Terms in Atomic Physics

Atoms and Subatomic Particles

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• Atoms are the fundamental building blocks of matter consisting of a dense nucleus surrounded by electrons
• Subatomic particles include protons, neutrons, and electrons
  • Protons and neutrons make up the atomic nucleus while electrons orbit the nucleus
• The number of protons in an atom's nucleus determines its atomic number and defines the element
• The mass number of an atom is the sum of the number of protons and neutrons in its nucleus

Molecules and Chemical Bonding

• Molecules are formed when two or more atoms bond together through chemical interactions
  • For example, water ($\text{H}_2\text{O}$) is a molecule consisting of two hydrogen atoms bonded to one oxygen atom
• Chemical bonding occurs when atoms share or transfer electrons to achieve a more stable electronic configuration
  • Covalent bonds involve the sharing of electrons between atoms (carbon dioxide, $\text{CO}_2$)
  • Ionic bonds involve the transfer of electrons from one atom to another, creating ions that attract each other (sodium chloride, $\text{NaCl}$)

Atomic Mass and the Periodic Table

Atomic Mass and Isotopes

• Atomic mass is the average mass of an element's atoms, typically expressed in atomic mass units (amu)
• The atomic mass of an element is a weighted average based on the relative abundances of its naturally occurring isotopes
  • For example, carbon has two stable isotopes: carbon-12 ($^{12}\text{C}$) and carbon-13 ($^{13}\text{C}$), with abundances of 98.93% and 1.07%, respectively
• Isotopes are atoms of the same element with different numbers of neutrons, resulting in varying mass numbers but the same atomic number and chemical properties
  • Isotopes have different nuclear stabilities and radioactive properties (uranium-235 and uranium-238)

Periodic Table Organization

• In the periodic table, elements are arranged in order of increasing atomic number, with atomic mass generally increasing from left to right
• The periodic table organizes elements with similar properties into groups (columns) and periods (rows) based on their electron configurations
  • Group 1 elements (alkali metals) have one valence electron and are highly reactive (sodium, potassium)
  • Group 18 elements (noble gases) have full valence electron shells and are generally unreactive (neon, argon)
• Trends in properties such as atomic radius, ionization energy, and electronegativity can be observed across the periodic table

Isotopes, Ions, and Neutral Atoms

Neutral Atoms and Ions

• Neutral atoms have an equal number of protons and electrons, resulting in a net charge of zero
• Ions are atoms or molecules that have gained or lost electrons, resulting in a net positive (cation) or negative (anion) charge
  • Cations are formed when an atom loses one or more electrons (sodium ion, $\text{Na}^+$)
  • Anions are formed when an atom gains one or more electrons (chloride ion, $\text{Cl}^-$)
• The number of protons remains constant in an atom, regardless of its ionic state or isotopic composition

Isotopes and Mass Number

• Isotopes are atoms of the same element with different numbers of neutrons, resulting in varying mass numbers but the same atomic number and chemical properties
  • For example, hydrogen has three isotopes: protium ($^1\text{H}$), deuterium ($^2\text{H}$), and tritium ($^3\text{H}$)
• The mass number of an isotope is the sum of the number of protons and neutrons in its nucleus
  • Carbon-12 has 6 protons and 6 neutrons, resulting in a mass number of 12
  • Carbon-13 has 6 protons and 7 neutrons, resulting in a mass number of 13

Atomic Notation for Elements and Isotopes

Standard Atomic Notation

• Atomic notation is written as $^A_Z\text{X}$, where X is the chemical symbol, Z is the atomic number (number of protons), and A is the mass number (protons + neutrons)
• The atomic number (Z) is written as a subscript to the left of the chemical symbol, while the mass number (A) is written as a superscript to the left of the chemical symbol
  • For example, carbon-12 is represented as $^{12}_6\text{C}$, indicating an atom with 6 protons, 6 neutrons, and 6 electrons
• Neutral atoms have the same number of protons and electrons, so the atomic number also represents the number of electrons in a neutral atom

Isotope Notation

• Isotopes of an element have the same atomic number but different mass numbers
  • Carbon-12 is represented as $^{12}_6\text{C}$, while carbon-13 is represented as $^{13}_6\text{C}$
• The mass number (A) distinguishes between isotopes of the same element
  • Uranium-235 ($^{235}_{92}\text{U}$) and uranium-238 ($^{238}_{92}\text{U}$) are isotopes of uranium with different numbers of neutrons
• Isotope notation is crucial for understanding nuclear reactions and radioactive decay processes