4.1

Electronic Configurations and Periodic Trends

Introduction

  • Focus of Module 3: Electronic configurations and periodic trends.

  • Review of quantum numbers: Four quantum numbers are essential for understanding electron locations around the nucleus:

    • Principal Quantum Number (N): Indicates the energy level of the electron.

    • Angular Momentum Quantum Number (L): Indicates the shape of the orbital (s, p, d, f).

    • Magnetic Quantum Number (M_L): Specifies the orbital orientation.

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

Electronic Configurations

  • Definition: Describes the arrangement of electrons in an atom.

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

    • Example with Helium:

    • Atomic number: 2 (2 electrons).

    • Both electrons in the 1s orbital have different spin quantum numbers:

      • First electron: (n=1, l=0, ML=0, Ms=+1/2)

      • Second electron: (n=1, l=0, ML=0, Ms=-1/2)

  • Implementation of the Pauli Exclusion Principle: Each orbital can hold a maximum of 2 electrons with opposite spins.

Aufbau Principle

  • Definition: Electrons fill orbitals starting from the lowest energy level to the highest.

  • Periodic Table as a Guide:

    • Hydrogen to Francium: Each left to right movement in the periodic table corresponds to adding one proton and one electron.

    • Example:

    • Hydrogen (1s¹): 1 electron

    • Helium (1s²): 2 electrons

    • Lithium (2s²): 3 electrons, starts filling the next energy level

    • Repeat process across periods: 1s, 2s, 2p, 3s, 3p, etc.

  • Group of Electrons:

    • 1s² (Helium)

    • 2s² (Beryllium)

    • 2p⁶ (Neon)

Energy Levels Effect on Configurations

  • Ordering of Orbitals: Electrons ordered by increasing energy:

    • 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.

  • Filling Order:

    • 1s² ➔ 2s² ➔ 2p⁶ ➔ 3s² ➔ 3p⁶ ➔ 4s² ➔ 3d¹⁰ ➔ 4p⁶, etc.

Hund's Rule

  • Definition: In degenerate orbitals, electrons fill each orbital singly first before pairing occurs.

    • Example with Carbon (Atomic Number 6):

    • 1s² 2s² 2p²:

    • Orbital diagram must represent Hund’s Rule: one electron in each 2p orbital before pairing could occur.

Orbital and Electron Configuration Examples

  • Carbon:

    • Configuration: 1s² 2s² 2p²

  • Sulfur (Atomic Number 16):

    • Configuration: 1s² 2s² 2p⁶ 3s² 3p⁴

  • Bromine (Atomic Number 35) using Noble Gas Shorthand:

    • Full configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁵

    • Noble Gas shorthand: [Ar] 4s² 3d¹⁰ 4p⁵.

Exceptions to Electronic Configurations

  • Notable Exceptions:

    • Chromium (Cr): Expected configuration would be [Ar] 4s² 3d⁴; instead, the actual one is [Ar] 4s¹ 3d⁵ (stability from half-full d subshell).

    • Copper (Cu): Expected: [Ar] 4s² 3d⁹; instead, actual is [Ar] 4s¹ 3d¹⁰ (stability of fully filled d subshell).

Applying Configurations to Ions

  • For ions: The electronic configurations still follow the same principles:

    • Chlorine (Cl):

    • Configuration: [Ne] 3s² 3p⁵

    • Chloride Ion (Cl⁻): [Ne] 3s² 3p⁶

    • Sodium Ion (Na⁺):

    • Configuration: [Ne] 3s¹

    • Na⁺: [Ne] (removing from the highest energy occupied orbital, which is 3s).

Transition Metals Specifics

  • When removing electrons from transition metals, the order generally stays constant where electrons are removed from the outermost shell (4s) before the inner d shell (3d).

  • General Configuration for Transition Metals:

    • Notation: n^2 (4s) and (n−1)d.

    • Example with Vanadium:

    • Neutral V: [Ar] 4s² 3d³

    • Vanadium Ion (V⁺): Electron removed from 4s: [Ar] 3d³

Magnetic Properties of Elements

  • Diamagnetic: Atoms with all paired electrons, essentially non-polarizable, not attracted to magnets.

  • Paramagnetic: Atoms with unpaired electrons that respond to magnetic fields.

    • Example:

    • Zinc: [Ar] 4s² 3d¹⁰, all paired ➔ Diamagnetic.

    • Copper: [Ar] 4s¹ 3d¹⁰, unpaired electrons ➔ Paramagnetic.

Conclusion

  • The understanding of electronic configurations and periodic trends is foundational for predicting element behavior, electron arrangement, and chemical reactions. Following the rules and exceptions noted is crucial for accurate representations.