Study Notes on Transition Metals & Coordination Chemistry

Definitions and Classification

• Transition Metals: Elements with partly filled d-orbitals in at least one common oxidation state. Group 12 elements are excluded due to fully filled d-orbitals.

• Scope: Focuses on comparing 1st, 2nd, and 3rd row d-block elements.

Key Trends Across Periods and Groups

• Atomic and Ionic Radii:

  • Decreases across a period due to increased effective nuclear charge ($Z_{eff}$).

  • Increases down a group, though the 2nd and 3rd rows have similar radii due to the Lanthanide Contraction (poor shielding by f-orbitals).

• Ionisation Energies (IE):

  • Generally increases across periods with minor dips at half-filled configurations.

  • Lanthanide contraction causes higher IEs in the 3rd row compared to the 2nd row.

• Melting Points: Peaks mid-series due to maximal unpaired electrons for M-M bonding, then decreases as orbitals fill. Generally higher in heavier transition metals.

Oxidation States

• Variability: Arises from the small energy gap between $ns$ and $(n-1)d$ orbitals.

• Trends:

  • 1st row commonly exhibits +2 and +3 states.

  • 2nd and 3rd rows favor higher oxidation states (e.g., $Os$ reaching +8) due to larger atomic size and lower relative attraction to valence electrons.

Coordination Chemistry and Bonding

• Bonding Character:

  • 1st row complexes are predominantly ionic and described by Crystal Field Theory (CFT).

  • 2nd and 3rd row complexes exhibit greater covalency due to better orbital overlap.

• Magnetic Properties:

  • 1st row often forms high-spin, paramagnetic complexes.

  • 2nd and 3rd rows typically form low-spin complexes due to larger crystal field splitting (\Delta).

• Geometry and Color:

  • Colors result from d-d transitions and charge transfers (MLCT and LMCT).

  • Heavier metals accommodate higher coordination numbers and diverse geometries like square planar or octahedral.