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
The concept of atoms dates back to ancient Greece but became scientific with John Dalton in 1805, who described atoms as indivisible particles unique to each element.
J.J. Thomson discovered the electron via the cathode ray tube experiment in 1897, identifying atoms as divisible and containing negative particles.
Millikan’s oil drop experiment measured the charge and calculated the mass of the electron (9.11 × 10⁻²⁸ g).
Rutherford’s gold foil experiment revealed the existence of a small, dense, positively charged nucleus surrounded by mostly empty space.
Chadwick discovered the neutron in 1932, completing the basic model of the atom: nucleus with protons and neutrons, surrounded by electrons.
The Bohr Model and Atomic Structure
Bohr proposed electrons orbit the nucleus in fixed energy levels (shells), and energy is required to move between them.
The maximum number of electrons in a shell is determined by the formula 2n².
Atomic number = number of protons; Mass number = protons + neutrons.
Isotopes are atoms of the same element with different numbers of neutrons and thus different mass numbers.
Average atomic mass is calculated based on the relative abundance of isotopes.
Valence Electrons and Lewis Structures
Valence electrons are the electrons in the outermost shell and determine chemical behavior.
Lewis structures represent valence electrons as dots around the chemical symbol.
Atoms gain, lose, or share electrons to complete their outer shell, forming ions or bonds.
Atomic Spectra and Energy Levels
According to Bohr, electrons absorb energy to jump to higher energy levels (excited state) and release energy when returning to lower levels (ground state).
The emitted energy appears as discrete light lines in an atomic spectrum, each unique to an element.
Series of transitions include the Lyman series (UV, to n=1), Balmer series (visible, to n=2), and Paschen series(infrared, to n=3).
Spectroscopy uses this principle to identify elements based on their emission lines.
The Wave-Mechanical Model and Orbitals
Bohr’s model couldn’t explain atoms with multiple electrons. The wave-mechanical model replaced it using quantum mechanics.
De Broglie proposed that particles like electrons also behave as waves.
Schrödinger developed equations describing orbitals—3D probability regions where electrons are likely found.
Heisenberg’s Uncertainty Principle states that we cannot know both the position and velocity of an electron simultaneously.
Quantum Numbers and Orbital Shapes
Each electron is defined by four quantum numbers:
Principal (n): energy level (1, 2, 3…)
Angular momentum (ℓ): shape (s = sphere, p = dumbbell, d = clover)
Magnetic (mₗ): orientation in space
Spin (mₛ): +½ or −½
Pauli Exclusion Principle: No two electrons can have the same four quantum numbers.
s orbitals hold 2 electrons, p orbitals 6, d orbitals 10, and f orbitals 14.
Rules for Electron Configuration
Aufbau Principle: Electrons fill the lowest available energy levels first.
Hund’s Rule: Electrons fill orbitals singly before pairing.
Not all orbitals fill in strict numerical order (e.g., 4s fills before 3d).
Electron configurations and noble gas notation show how electrons are arranged. For example: Na = [Ne] 3s¹.
Transition Elements and Electron Behavior
Transition metals have partially filled d orbitals, leading to:
Multiple oxidation states
Colored compounds
Catalytic activity
Complex ion formation
Some elements (like Cr and Cu) have exceptions to predicted configurations for added stability.
The Periodic Table and Its Trends
Mendeleev created the first Periodic Table based on atomic weights and recurring properties.
Moseley later revised it by ordering elements by atomic number, forming the modern Periodic Law.
Periods are horizontal rows (1–7), and groups are vertical columns (1–18).
Metals are on the left (most active: bottom-left), nonmetals on the right (most active: top-right), and metalloidslie along the stair-step line.
Periodic Trends to Know
Atomic radius:
Decreases across a period (left to right)
Increases down a group (top to bottom)
Ionic radius:
Cations (metals): smaller than atomic radius
Anions (nonmetals): larger than atomic radius
Electronegativity:
Increases across a period
Decreases down a group
Highest: Fluorine (F), lowest: Francium (Fr)
Ionization energy:
Energy required to remove an electron
Increases across a period, decreases down a group
Peaks at noble gases due to stable configurations
Radioactivity and Nuclear Decay
Types of radiation:
Alpha particles (α): Helium nuclei, low penetration, reduce atomic number by 2 and mass number by 4
Beta particles (β): High-speed electrons, moderate penetration, convert neutrons to protons
Gamma rays (γ): High-energy light, deep penetration, no mass or charge
Detection methods:
Photographic plates
Scintillation counters
Geiger counters
Half-Life and Radioactive Dating
Half-life is the time for half of a radioactive sample to decay.
Carbon-14 dating uses this concept to estimate the age of organic materials, like ancient wood or relics.
Nuclear Reactions
Fission: Splitting heavy nuclei (e.g., U-235) into smaller parts, releasing energy and neutrons—used in nuclear power and bombs.
Fusion: Combining light nuclei (e.g., H + H → He) into heavier ones—releases more energy than fission but requires extreme conditions.
Both processes convert mass into energy (E = mc²).
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.