Inorganic Chemistry Detailed Notes

Inorganic Chemistry Overview

  • Topic: Forensic Investigation in inorganic chemistry
  • Speaker: Dr. Mark D. Spicer, Department of Applied Science, GCU

1. Transition Elements and the Periodic Table

  • Properties of transition elements vary with their atomic number (Z).
    • Atomic number increases from left to right and top to bottom.

2. Electron Configurations

  • Fundamentals:
    • Essential for understanding chemical behavior.
    • The Aufbau Principle dictates the order of orbital filling:
    • 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p.

2.1 Core and Valence Electrons

  • Core Electrons: Not involved in bonding; equivalent to noble gas electron configurations (e.g., [Ar] for first-row transition elements).
  • Valence Electrons: Involved in bonding; for first transition series, they determine chemical properties.
  • Properties of example element Scandium (Z = 21):
    • Electron configuration: 1s^22s^22p^63s^23p^6 4s^23d^1
    • Core Electrons: 18, equivalent to Argon, valence electrons = 3.

3. Quantum Numbers and Orbitals

  • Four Quantum Numbers:
    • n (principal quantum number): Determines electron shell.
    • ext{ℓ} (azimuthal quantum number): Defines orbital shape.
    • m_ ext{ℓ} (magnetic quantum number): Determines orbital orientation.
    • m_s (spin quantum number): Indicates spin direction.
  • Pauli Exclusion Principle: Each electron must have a unique set of quantum numbers.

3.1 Orbital Shapes

  • s-orbitals: Spherical
  • p-orbitals: Dumbbell-shaped
  • d-orbitals: More complex shapes

4. Trends in Periodic Properties

4.1 Atomic Radii

  • Definition: Measure of the atom's size, typically from nucleus to electron shell boundary.
    • Varieties in measurement: Van der Waals radius, covalent radius, ionic radius.
    • General trends in radii across periods include:
    • Decreases across a period due to increasing nuclear charge.
    • Increases from period 4 to 5/6 due to inner electron shielding.

4.2 Effective Nuclear Charge (Z*)

  • Definition: The net positive charge experienced by outer electrons, expressed as: Z^* = Z - S
    • Increases across a period leading to decreased atomic radii.
    • Example calculation for Vanadium (Z=23):
    • S = (18 x 1.0) + (2 x 0) + (3 x 0.35) = 19.05
    • Z^* = 23 - 19.05 = 3.95

4.3 Ionization Energies

  • Definition: Energy needed to remove electrons.
    • Increases across periods; more energy needed due to decreasing atomic radii and increased nuclear charge.

5. Oxidation States

  • Definition: The charge of an ion post ionization.
    • Trends:
    • Transition elements frequently exhibit multiple oxidation states due to d-orbital involvement.
    • Stability varies through periodic table, with +2 oxidation states common for all d-elements.
    • Higher oxidation states favored at the beginning of transition period.
    • Example oxidation state stability:
      • Sc: +2, Cr: +3-6, Cu: +2, etc.

5.1 Born-Haber Cycle in Oxidation States

  • Describes lattice formation and energies needed to form compounds from elements.
    • Uses Hess’s Law for calculating enthalpy changes through individual steps.

5.2 High Oxidation States

  • Stabilized by bonded ligands (oxides, halides), which provide electron density and enhance covalency.

Conclusion

  • Understanding periodic properties, electron configurations, ionization states, and oxidation trends is essential in inorganic chemistry and forensic applications.

  • Prepared by Dr. Mark D. Spicer for Inorganic Chemistry: Trimester B 2020-21.