In-Depth Notes on Atomic Structure and Atomic Mass

Atoms consist of a nucleus surrounded by electrons.
Nucleus contains:
- Protons: positively charged particles.
- Neutrons: neutral particles.
Electrons are negatively charged and are held by electrostatic attraction to the nucleus.
The nucleus accounts for 99.97% of the atomic mass, though the electron cloud occupies a much larger volume (10³ - 10⁵ times larger than the nucleus).
Protons and neutrons are collectively termed nucleons.

Atomic Number (Z): Number of protons in the nucleus; determines the element.
Mass Number (A): Sum of protons and neutrons in the nucleus; indicates total mass of the nucleus.
- Formula: Number of Neutrons = A - Z.
Electrically Neutral Atoms: Equal numbers of protons and electrons.

Isotopes are atoms with the same atomic number but different mass numbers (A).
They have identical chemical properties but exhibit different physical properties (e.g., mass, density).
Examples:
- Carbon isotopes: Carbon-12, Carbon-13, Carbon-14.

Radioactive Isotopes (Radioisotopes): Isotopes that are unstable and emit radiation.
- All nuclei with 84 or more protons are radioactive.
Types of Radiation:
- Alpha (α) Particles: Heavy, positively charged (2 protons & 2 neutrons); low penetrating power (stopped by paper).
- Beta (β) Particles: High-energy electrons; negatively charged; more penetrating than alpha particles (increases atomic number by 1).
- Gamma (γ) Radiation: High-energy electromagnetic radiation; carries no charge; highly penetrating.

The band of stability indicates isotopes that are stable compared to those that are radioactive.
Unstable nuclei undergo decay processes emitting alpha, beta, or gamma radiation.

Stability is affected by the ratio of neutrons to protons (n:p).
Lighter nuclei have a balanced n:p ratio; heavy nuclei require higher neutrons to protons for stability.
Nuclei outside stability tend to emit particles to reach a stable n:p ratio.

Half-life: Time required for half of the atoms in a sample to undergo radioactive decay; it remains constant regardless of the sample size.

Relative Atomic Mass (Ar): Average mass of an atom accounting for isotopic abundance, compared to carbon-12.
Relative Isotopic Mass (Ir): Mass of an individual isotope.
Abundance of isotopes is measured using a mass spectrometer, which reveals the mixture and average mass from isotopes present.

Electrons are arranged in energy levels/shells around the nucleus, denoted as 1s, 2s, 2p, etc.
Electron configurations demonstrate how electrons occupy orbitals:
- Lower energy levels fill first.
- Maximum number of electrons per shell is determined by the formula 2n^2.
- For example, Hydrogen: 1s¹, Helium: 1s², Lithium: 1s² 2s¹, etc.
Valence Electrons: Electrons in the outermost shell; involved in chemical bonding and reactions following the octet rule.

The Bohr model explains hydrogen well but cannot account for more complex atoms nor why electron shells are filled the way they are.
Schrödinger's Model: Introduced quantized wave functions, describing the probability of finding electrons in orbitals instead of fixed paths.
Utilizes the concept of orbitals (s, p, d, f) where electron density varies, differs from classical orbits.
Pauli’s Exclusion Principle states that each orbital can hold a maximum of 2 electrons with opposite spins.
Anomalous Electron Configurations: Some transition metals (Cr, Cu) show unexpected electron arrangements for stability.

Understanding atomic structure includes knowledge of atomic particles, isotopes, radioactivity, electronic configurations, and quantum mechanics.
Different elements exhibit unique atomic behaviors due to the arrangement and number of their electrons, neutrons, and protons.