🟣 This is Higher Level (HL) content.
📘 d-Orbital Splitting
In an isolated gaseous transition metal ion, the five d-orbitals are degenerate (equal energy). When ligands bond to the metal, their lone pairs repel the d-electrons unequally, causing the d-orbitals to split into two energy levels.
The energy gap between the two levels is called ΔE
d-Orbital Splitting in an Octahedral Field
How Colour Arises
- White light passes through the solution
- A d-electron absorbs a photon with the exact energy matching ΔE
- The electron is promoted from the lower to the upper d-orbital level
- The remaining wavelengths are transmitted — the observed colour is the complementary colour of the absorbed light
The Colour Wheel
Factors Affecting ΔE and Colour
| Factor | Effect on ΔE | Effect on Colour |
|---|---|---|
| Nature of ligand | Strong field ligands (CN⁻) → large ΔE; Weak field (Cl⁻, H₂O) → small ΔE | Different ligands cause different absorption wavelengths |
| Oxidation state | Higher charge → ligands pulled closer → larger ΔE | Colour changes with oxidation state |
| Identity of metal | Different nuclear charges and configs → unique splitting | Each metal produces characteristic colours |
| Coordination geometry | Octahedral and tetrahedral fields split differently | Same metal/ligand but different geometry → different colour |
⚠️ Why Some are Colourless
Ions with d⁰ (e.g. Sc³⁺) or d¹⁰ (e.g. Zn²⁺, Cu⁺) configurations cannot undergo d-d transitions because there are either no electrons to promote or no empty higher-energy orbitals. These ions form colourless compounds.
⚠️ Exam Tip
If asked "why is a complex coloured?", always mention: (1) ligands cause d-orbital splitting, (2) d-electron absorbs visible light matching ΔE, (3) observed colour is the complementary colour. Use the equation ΔE = hf.