Inorganic Chemistry: Coordination and Transition Metal Chemistry
Core concepts
Master the vocabulary of coordination compounds: central metal, ligand, coordination number, oxidation state, and the difference between monodentate and polydentate ligands (EDTA is hexadentate). Apply IUPAC naming rules: ligands alphabetically, anionic ligands ending in -o, then the metal with its oxidation state in Roman numerals. Understand isomerism - geometrical (cis/trans in square planar and octahedral), optical, ionisation, hydrate and linkage isomerism.
Explain bonding using crystal field theory: d orbitals split into t₂g and e_g sets in an octahedral field, with Δ₀ determined by ligand strength (the spectrochemical series: I⁻ < Br⁻ < Cl⁻ < F⁻ < OH⁻ < H₂O < NH₃ < en < NO₂⁻ < CN⁻ < CO). Strong-field ligands give low-spin complexes; weak-field ligands give high-spin complexes. Use the magnetic moment μ = √(n(n+2)) Bohr magnetons to deduce unpaired electrons. Colour arises from d–d transitions; [Cu(NH₃)₄]²⁺ is deep blue while [Cu(H₂O)₆]²⁺ is pale blue. Know the chelate effect and the stability of inner-orbital versus outer-orbital complexes.
Traps to watch
- Counting the metal's oxidation state from the complex charge incorrectly.
- Forgetting that CN⁻ is strong-field while Cl⁻ is weak-field.
- Assuming all tetrahedral complexes are high spin (they essentially are, because Δ_t is small).
- Mixing up linkage isomerism (ambidentate ligands such as SCN⁻) with ionisation isomerism.
Worked example
[Fe(CN)₆]⁴⁻ contains Fe²⁺, a d⁶ ion. CN⁻ is strong field, so electrons pair up: t₂g⁶, zero unpaired electrons, diamagnetic and low spin. Replacing CN⁻ with H₂O gives a high-spin d⁶ complex with four unpaired electrons. Rule of thumb: count d electrons, identify ligand field strength, then fill the t₂g/e_g diagram.