Bonding: VBT and CFT

The essentials

Valence bond theory (VBT) says the metal uses hybridised orbitals to accept electron pairs from ligands. Six-coordinate complexes use d²sp³ (inner orbital, using (n−1)d) or sp³d² (outer orbital, using nd) hybrids and are octahedral; four-coordinate ones use sp³ (tetrahedral) or dsp² (square planar). Crystal field theory (CFT) treats ligands as point charges that split the d orbitals.

Key points

• [Co(NH₃)₆]³⁺: Co³⁺ d⁶, d²sp³, all paired, diamagnetic. [CoF₆]³⁻: sp³d², 4 unpaired electrons, paramagnetic.
• [NiCl₄]²⁻: sp³, tetrahedral, 2 unpaired (paramagnetic). [Ni(CN)₄]²⁻: dsp², square planar, diamagnetic.
• In an octahedral field, d orbitals split into lower t₂g (three) and higher e_g (two); the gap is Δo. In a tetrahedral field the order is reversed and Δt = (4/9)Δo, so tetrahedral complexes are almost always high spin.
• If Δo < P (pairing energy), the fourth electron enters e_g (weak field, high spin, t₂g³e_g¹). If Δo > P, it pairs in t₂g (strong field, low spin, t₂g⁴e_g⁰).
• Spectrochemical series (increasing field strength): I⁻ < Br⁻ < SCN⁻ < Cl⁻ < S²⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < edta⁴⁻ < NH₃ < en < CN⁻ < CO.
• Colour comes from d–d transitions; the observed colour is complementary to the absorbed light. [Ti(H₂O)₆]³⁺ (d¹) is violet. Anhydrous CuSO₄ is white, hydrated CuSO₄ is blue.
• Limitations: VBT does not explain colour or distinguish weak and strong ligands well; CFT treats ligands as point charges and cannot explain why some anions are weak-field ligands.

Exam tips

• Steps: find oxidation state → d count → ligand strength → electron arrangement → unpaired electrons → μ.

Common mistakes

• Assuming all Ni(II) four-coordinate complexes are tetrahedral; strong-field CN⁻ gives square planar.