Unit 1: Comprehensive Summary

Atoms are the fundamental units of matter, composed of protons (positive charge, ~1 amu mass), neutrons (no charge, ~1 amu mass) in the nucleus, and electrons (negative charge, negligible mass) orbiting in shells. The atomic number (Z) defines an element by its number of protons. The mass number (A) is the sum of protons and neutrons.

Isotopes are atoms of the same element (same Z) but with different numbers of neutrons (different A). They have identical chemical properties due to the same electron configuration.

Ions are charged atoms formed by gaining or losing electrons. Cations are positive (lose electrons), and anions are negative (gain electrons).

The relative atomic mass ($A_r$) is the weighted average of the masses of an element's naturally occurring isotopes, relative to 1/12th the mass of a Carbon-12 atom.

Electrons occupy discrete energy levels (shells: K, L, M, ... or $n=1, 2, 3, \dots$), which are further divided into sub-levels (s, p, d, f). Each sub-level contains specific orbitals: 1 s-orbital (spherical), 3 p-orbitals (dumbbell-shaped), 5 d-orbitals, and 7 f-orbitals.

The arrangement of electrons is governed by three rules:

• Pauli Exclusion Principle: An orbital holds a maximum of two electrons with opposite spins.
• Aufbau Principle: Electrons fill orbitals from lowest to highest energy (e.g., 1s, 2s, 2p, 3s, 3p, 4s, 3d...).
• Hund's Rule: For degenerate orbitals, electrons fill singly with parallel spins before pairing up.

Electronic configurations can be written in orbital notation (e.g., $1s^2 2s^2 2p^4$) or shorthand using noble gas symbols (e.g., $[He] 2s^2 2p^4$). For ions, electrons are added or removed, typically from the highest principal energy level first for cations (e.g., 4s before 3d in transition metals).

The Periodic Table arranges elements by increasing atomic number, revealing periodicity in their properties, influenced by nuclear charge, shielding, and electron repulsion.

• Atomic Size: Decreases across a period (increased nuclear charge pulls electrons closer); Increases down a group (more electron shells, increased shielding). Cations are smaller than parent atoms; anions are larger.
• First Ionization Energy ($IE_1$): Energy to remove the first electron. Generally increases across a period (stronger nuclear attraction); Decreases down a group (increased size and shielding). Exceptions occur due to subshell filling and electron-electron repulsion (e.g., Mg > Al, N > O).
• First Electron Affinity ($EA_1$): Energy change when an electron is added. Generally becomes more negative across a period (stronger attraction for incoming electron); Becomes less negative down a group (increased size and shielding). Exceptions also occur due to subshell filling and repulsion.
• Electronegativity: Tendency of an atom to attract bonding electrons. Increases across a period; Decreases down a group. Fluorine is the most electronegative.
• Metallic Character: Ease of losing electrons. Decreases across a period (metals to nonmetals); Increases down a group. Nonmetallic character shows the opposite trend.

Our understanding of the atom has evolved through scientific inquiry:

• Democritus (Ancient Greece): Proposed indivisible particles called "atomos." Philosophical, not experimental.
• Dalton's Atomic Theory (1803): Atoms are indivisible, identical for an element, combine in whole-number ratios. Explained laws of conservation of mass and definite/multiple proportions.
  • Modifications: Atoms are divisible (subatomic particles), isotopes exist (atoms of same element not identical in mass), atoms can be transformed (nuclear reactions).
• Thomson's Plum Pudding Model (1897): Discovered the electron using cathode ray tubes. Proposed a sphere of positive charge with embedded electrons.
• Rutherford's Nuclear Model (1909): Gold foil experiment showed most alpha particles passed through, but some were deflected. Concluded atom is mostly empty space with a tiny, dense, positively charged nucleus and electrons orbiting it. Flaw: classical physics predicted electron spiral.
• Bohr's Model (1913): Explained hydrogen's line spectrum. Proposed quantized energy levels where electrons orbit without radiating energy. Electrons move between levels by absorbing/emitting photons. Limited to one-electron species.
• Quantum Mechanical Model (1926 - Present): Developed by Schrödinger and others. Electrons exist in atomic orbitals (regions of probability), not fixed orbits. Incorporates wave-particle duality and Heisenberg Uncertainty Principle. Electrons described by four quantum numbers ($n, l, m_l, m_s$). Most accurate model, explaining multi-electron atoms and bonding.