General Chemistry Study Notes: Atomic Structure, Electron Configurations, and Periodic Trends

Atomic Structure, Electron Configurations, and Periodic Trends

Atomic Structure

The structure of the atom is fundamental to understanding chemical properties and behavior. Atoms consist of a nucleus containing protons and neutrons, surrounded by electrons in defined energy levels.

• Protons: Positively charged particles found in the nucleus.

• Neutrons: Neutral particles found in the nucleus.

• Electrons: Negatively charged particles found in orbitals around the nucleus.

• Atomic Number (Z): Number of protons in the nucleus; defines the element.

• Mass Number (A): Sum of protons and neutrons in the nucleus.

• Isotopes: Atoms of the same element with different numbers of neutrons.

Electron Configurations

Electron configuration describes the arrangement of electrons in an atom's orbitals. The distribution follows the Aufbau principle, Pauli exclusion principle, and Hund's rule.

• Aufbau Principle: Electrons fill the lowest energy orbitals first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.

• Hund's Rule: Electrons occupy orbitals singly before pairing up.

• Notation: Electron configurations are written using subshell notation (e.g., 1s2 2s2 2p6).

Example: The electron configuration of oxygen (O) is 1s2 2s2 2p4.

Quantum Numbers

Quantum numbers describe the properties of atomic orbitals and the electrons in them.

• Principal Quantum Number (n): Indicates the main energy level (shell).

• Angular Momentum Quantum Number (l): Indicates the subshell (s, p, d, f).

• Magnetic Quantum Number (ml): Indicates the orientation of the orbital.

• Spin Quantum Number (ms): Indicates the spin direction of the electron (+1/2 or -1/2).

Periodic Trends

The periodic table organizes elements by increasing atomic number and reveals trends in properties such as atomic radius, ionization energy, and electronegativity.

• Atomic Radius: Decreases across a period (left to right), increases down a group.

• Ionization Energy: Increases across a period, decreases down a group.

• Electronegativity: Increases across a period, decreases down a group.

• Electron Affinity: Generally becomes more negative across a period.

Example: Fluorine has the highest electronegativity in the periodic table.

Important Equations

• Energy of a photon: $E = h \nu$ Where: E = energy, h = Planck's constant, \nu = frequency

• Relationship between wavelength and frequency: $c = \lambda \nu$ Where: c = speed of light, \lambda = wavelength, \nu = frequency

Periodic Table Classification

The periodic table is divided into groups (columns) and periods (rows). Elements are classified as metals, nonmetals, and metalloids.

• Groups: Elements with similar chemical properties (e.g., alkali metals, halogens).

• Periods: Elements with the same number of electron shells.

• Metals: Good conductors, malleable, ductile.

• Nonmetals: Poor conductors, brittle.

• Metalloids: Properties intermediate between metals and nonmetals.

Sample Table: Periodic Trends

The following table summarizes key periodic trends:

Property	Across Period (Left to Right)	Down Group (Top to Bottom)
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electronegativity	Increases	Decreases
Electron Affinity	Becomes more negative	Becomes less negative

Orbital Diagrams

Orbital diagrams visually represent the arrangement of electrons in orbitals, showing the application of Hund's rule and the Pauli exclusion principle.

Additional info: These notes expand on the brief points in the original material, providing definitions, examples, and context for general chemistry students.