Lecturers and Consultation

  • Dr PS Ramaripa, Office: N1010, EXT.: 3783
  • Mr BN Mbokane, Office: N3019, EXT.: 4771
  • Consultation Times: Mon – Thurs, 09:00 – 15:00

Key Dates

  • Test 1: 14/04
  • Test 2: 05/05
  • Assignment 1: 01/04
  • Assignment 2: 30/04

Course Content

  1. Foundations of Inorganic Chemistry
    • Atomic Structure
    • Effective Nuclear Charge (ENC): Slater’s Rules
    • Quantum Chemistry
    • Periodic Properties
    • Chemical Bonding (Lewis, VBT, MOT)
  2. Chemical Bonding, Elements, and Compounds
    • Solid State Chemistry
    • Transition Metals
    • Coordination Chemistry
    • Isomers
  3. Acid-Base Chemistry and Electrochemistry
    • Acids & Bases
    • Oxidation & Reduction
    • Latimer, Frost, and Pourbaix Diagrams

Inorganic Chemistry Definition

  • Involves the chemistry of all elements in the Periodic Table including carbon.

Periodic Table

  • Developed by Dmitri Mendeleev.
  • Modern periodic table is organized by increasing atomic number, rather than atomic weight.
  • Periodic Law: Elements display regular patterns of properties when arranged by atomic number.
  • Period Number: Indicates the highest energy level of electrons in the element.
  • Groups: Elements share properties and have the same valence electron arrangement.

Effective Nuclear Charge (Z*)

  • Z∗=Z−SZ^* = Z - S, where Z = nuclear charge, S = shielding constant.
  • Z* decreases with electron shielding from inner electrons.

Electron Configuration Principles

  • Aufbau Principle: Fill lower energy orbitals first.
  • Hund’s Rule: Electrons fill degenerate orbitals singly before pairing.
  • Pauli Exclusion Principle: No two electrons in an atom can have the same quantum numbers.

Key Concepts in Atomic Structure

  • Atomic orbitals categorized into s, p, d, f types.
  • Electrons defined by unique quantum numbers: principal, subsidiary, magnetic, and spin.

Quantum Mechanics and Electrons

  • Heisenberg Uncertainty Principle: Position and momentum of particles can't be determined simultaneously.
  • Schrödinger Equation: Describes the wave function of electrons and allows energy levels to be quantified.

Periodic Trends

  • Atomic Size: Decreases left to right, increases top to bottom.
  • Ionization Energy: Increases left to right, decreases top to bottom.
  • Electron Affinity: Higher values for elements attracting electrons more readily.

Electronegativity

  • Defined as the ability of an atom to attract electrons in a bond.
  • Influenced by atomic size and electron configuration.
  • Various scales: Pauling, Mulliken, Allred-Rochow.
  • Higher electronegativity correlates with smaller atomic radius and higher nuclear charge.