Electron Configuration and Periodic Trends: Study Notes

Quantum Mechanics, Electron Configurations, and Periodic Trends

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom's orbitals. Understanding the order in which electrons fill these orbitals is essential for predicting chemical properties and reactivity.

• Order of Filling (Aufbau Principle): Electrons fill orbitals in order of increasing energy, following the sequence: 1s, 2s, 2p, 3s, 3p, 4s before 3d, 4p, 5s before 4d, and so on. This is often remembered using the diagonal rule.

• Notation: Electron configurations are written using superscripts to indicate the number of electrons in each subshell. Example: 1s^2 2s^2 2p^4.

• Noble Gas Core Notation: To simplify configurations, use the symbol of the previous noble gas in brackets. Example: Chlorine is [Ne] 3s^2 3p^5.

• Exceptions: Some elements have electron configurations that deviate from the expected order due to increased stability of half-filled or fully-filled subshells. Notable examples:

  • Chromium: [Ar] 4s^1 3d^5

  • Copper: [Ar] 4s^1 3d^{10}

Example: The electron configuration of oxygen (atomic number 8) is 1s^2 2s^2 2p^4.

Orbital Box Diagrams and Hund's Rule

Orbital box diagrams visually represent the arrangement of electrons and their spins within orbitals. These diagrams help illustrate the application of Hund's Rule and the Pauli Exclusion Principle.

• Boxes: Each box represents an orbital, which can hold a maximum of two electrons.

• Arrows: An upward arrow (↑) indicates an electron with spin-up; a downward arrow (↓) indicates spin-down.

• Hund's Rule: When filling orbitals of equal energy (degenerate orbitals, such as the three 2p orbitals), place one electron in each orbital with parallel spins before pairing any electrons.

• Pauli Exclusion Principle: No two electrons in the same orbital can have the same set of quantum numbers; thus, they must have opposite spins.

Example: The 2p electrons of nitrogen (atomic number 7) are arranged as ↑ ↑ ↑ in three separate boxes before any pairing occurs.

Periodic Trends

Periodic trends describe how certain atomic properties change in a predictable way across periods (rows) and groups (columns) of the periodic table. Understanding these trends is crucial for predicting element behavior.

Atomic Radius

• Across a Period (→): Atomic radius decreases from left to right due to increasing nuclear charge, which pulls electrons closer to the nucleus.

• Down a Group (↓): Atomic radius increases because additional electron shells are added, making the atom larger.

Ionization Energy (IE)

• Definition: The energy required to remove an electron from a gaseous atom.

• Across a Period (→): Ionization energy increases as atoms become smaller and electrons are held more tightly.

• Down a Group (↓): Ionization energy decreases because the outermost electron is farther from the nucleus and more easily removed.

Electronegativity

• Definition: The tendency of an atom to attract electrons in a chemical bond.

• Across a Period (→): Electronegativity increases as atoms more strongly attract electrons.

• Down a Group (↓): Electronegativity decreases as atomic size increases and the nucleus has less pull on bonding electrons.

• Note: Noble gases are typically excluded from electronegativity trends because they rarely form bonds.

Ionic Radius

• Cations (+): Positively charged ions are smaller than their neutral atoms because they lose electrons, reducing electron-electron repulsion and sometimes losing an entire electron shell.

• Anions (−): Negatively charged ions are larger than their neutral atoms due to increased electron-electron repulsion after gaining electrons.

• Isoelectronic Ions: For ions with the same number of electrons, the ion with more protons (higher atomic number) is smaller because the electrons are pulled closer to the nucleus.

Trend	Across a Period (→)	Down a Group (↓)
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electronegativity	Increases	Decreases

Example: Sodium (Na) has a larger atomic radius than chlorine (Cl) in the same period. Fluorine (F) has a higher electronegativity than lithium (Li).

Additional info: These trends are explained by effective nuclear charge, electron shielding, and quantum mechanical principles governing electron arrangement.