Inorganic Chemistry and Coordination Compounds
Scope
This domain examines coordination compounds: nomenclature, coordination number and geometry, isomerism, crystal field splitting, high- versus low-spin configurations, magnetic moments, and the chelate effect.
Core rules
- Coordination number is the count of donor atoms bonded to the metal; oxidation state is found by subtracting ligand charges from the complex charge.
- Common geometries: CN 2 linear, CN 4 tetrahedral or square planar, CN 6 octahedral.
- Crystal field theory splits d orbitals: octahedral gives t2g (lower) and eg (upper) with Δo; tetrahedral splitting is smaller and inverted in order.
- Strong-field ligands (CN⁻, CO, phen) produce low-spin, large Δo; weak-field ligands (I⁻, Br⁻, F⁻, H₂O) produce high-spin, small Δo. The spectrochemical series orders them.
- Magnetic moment μ = √(n(n+2)) Bohr magnetons, where n is unpaired electrons.
Worked example
Compare [Fe(CN)₆]³⁻ and [FeF₆]³⁻. In both, Fe is +3, giving d⁵. Cyanide is a strong-field ligand: the d⁵ electrons pair into t2g⁵, leaving one unpaired electron, so μ = √(1×3) = 1.73 BM and the complex is low spin. Fluoride is weak field: the electrons spread across all five d orbitals as t2g³eg², leaving five unpaired electrons, so μ = √(5×7) = 5.92 BM and the complex is high spin.
Traps
- Counting d electrons without subtracting the metal's oxidation state from its group number.
- Assuming tetrahedral complexes are usually low spin; their smaller splitting favors high spin.
- Confusing coordination number with oxidation state or with the number of ligands when polydentate ligands bind.
- Missing the chelate effect, which makes multidentate complexes more thermodynamically stable.