Rucete ✏ SAT Chemistry In a Nutshell
2. Atomic Structure and the Periodic Table of the Elements
Understanding atoms is essential to mastering chemistry. This chapter covers the development of atomic theory, models of the atom, quantum numbers, electron configurations, and trends in the periodic table. These ideas explain how atoms behave and interact, and how we organize them into the Periodic Table.
The Evolution of Atomic Theory
- Dalton: Atoms are indivisible particles unique to each element.
- Thomson: Discovered electrons with cathode ray tube; atoms are divisible.
- Millikan: Measured charge and mass of electrons.
- Rutherford: Gold foil experiment revealed nucleus is dense and positive.
- Chadwick: Discovered the neutron in 1932.
The Bohr Model and Atomic Structure
- Electrons orbit the nucleus in fixed energy levels (2n² rule).
- Atomic number = protons; Mass number = protons + neutrons.
- Isotopes = same element, different neutrons → different mass.
- Average atomic mass = weighted average of isotopes.
Valence Electrons and Lewis Structures
- Valence electrons are outermost electrons, determine bonding.
- Lewis structures use dots around element symbols to show valence electrons.
- Atoms bond by gaining, losing, or sharing electrons to fill their shells.
Atomic Spectra and Energy Levels
- Electrons absorb energy to jump up (excited), release light when falling down (ground).
- Emission lines (atomic spectra) are unique to each element.
- Lyman (UV), Balmer (visible), Paschen (IR) series.
- Spectroscopy identifies elements using emission lines.
The Wave-Mechanical Model and Orbitals
- Wave-particle duality (de Broglie): electrons have wave behavior.
- Schrödinger: Orbitals = 3D probability regions.
- Heisenberg: Cannot know both position and velocity of an electron exactly.
Quantum Numbers and Orbital Shapes
- Principal (n): Energy level
- Angular momentum (ℓ): Shape (s, p, d, f)
- Magnetic (mₗ): Orientation
- Spin (mₛ): +½ or −½
- Pauli Exclusion: No two electrons share same 4 quantum numbers.
- s (2e⁻), p (6), d (10), f (14)
Rules for Electron Configuration
- Aufbau: Fill lowest energy first.
- Hund’s Rule: Fill orbitals singly before pairing.
- 4s fills before 3d, etc.
- Use noble gas shorthand: Na = [Ne] 3s¹
Transition Elements and Electron Behavior
- Partially filled d orbitals → multiple oxidation states, colored compounds, catalysts, complex ions.
- Cr, Cu: exceptions in configuration for stability.
The Periodic Table and Its Trends
- Mendeleev: First table based on atomic weight.
- Moseley: Modern table based on atomic number.
- Periods = rows; groups = columns.
- Metals: Left; Nonmetals: Right; Metalloids: Stair-step line.
Periodic Trends to Know
- Atomic radius: Decreases across period, increases down group.
- Ionic radius: Cations smaller, anions larger than atom.
- Electronegativity: Increases across, decreases down. Highest: F, Lowest: Fr.
- Ionization energy: Increases across, decreases down. Peaks at noble gases.
Radioactivity and Nuclear Decay
- Alpha (α): Low penetration, −2 atomic number, −4 mass.
- Beta (β): Moderate penetration, neutron → proton.
- Gamma (γ): High-energy light, deep penetration, no mass.
- Detection: photographic plates, scintillation counters, Geiger counters.
Half-Life and Radioactive Dating
- Half-life = time for half of substance to decay.
- Carbon-14 used for dating ancient biological materials.
Nuclear Reactions
- Fission: Heavy nuclei split → energy + neutrons (e.g., U-235).
- Fusion: Light nuclei combine → more energy (e.g., H + H → He).
- Both follow E = mc² (mass → energy).
In a Nutshell
This chapter explores the history and structure of atoms, quantum mechanics, periodic trends, electron configurations, and nuclear chemistry. From Dalton to Schrödinger, and from electron orbitals to radioactive decay, it builds the atomic-level understanding needed to explain chemical behavior and periodic organization.