Colour in Transition Metal Compounds
This lesson explores why transition metal compounds are coloured, the basics of crystal field theory, the factors that affect colour, and how colorimetry and the Beer–Lambert law are used to determine concentrations of coloured solutions (AQA 3.2.5).
Why Are Transition Metal Compounds Coloured?
The d-Orbital Splitting Model
In a free (gaseous) transition metal ion, the five 3d orbitals are degenerate (they all have the same energy). When ligands approach the metal ion to form a complex, the electrostatic interaction between the ligand lone pairs and the d electrons causes the d orbitals to split into two energy levels.
For an octahedral complex:
- Three d orbitals are lowered in energy (the t₂g set: dₓy, dₓz, dyz).
- Two d orbitals are raised in energy (the eg set: dₓ²₋y², dz²).
- The energy gap between them is called Δ (delta) or Δₒct.
For a tetrahedral complex:
- The splitting pattern is inverted (two lower, three higher).
- The splitting energy Δₜₑₜ is smaller than Δₒct (roughly 4/9 of Δₒct).
The d–d Transition
- When white light passes through a solution of a transition metal complex, photons of a specific frequency are absorbed.
- The energy of the absorbed photon corresponds to the energy gap Δ between the split d orbitals: ΔE = hν = hc/λ.
- An electron is promoted from the lower set of d orbitals to the upper set. This is called a d–d transition.
- The light that is not absorbed passes through and is perceived as the complementary colour of the absorbed light.
The Colour Wheel
| Colour absorbed | Colour observed (complementary) |
|---|
| Red | Green |
| Orange | Blue |
| Yellow | Violet |
| Green | Red |
| Blue | Orange |
| Violet | Yellow |
Example
[Cu(H₂O)₆]²⁺ absorbs light in the red/orange region of the visible spectrum. The complementary colour is blue, which is why copper(II) sulfate solution appears blue.
Factors Affecting Colour
Three factors change the size of the d-orbital splitting (Δ) and therefore change the colour of a complex:
1. The Ligand
Different ligands cause different amounts of splitting. This is summarised in the spectrochemical series (from weak field to strong field):
I⁻ < Br⁻ < Cl⁻ < F⁻ < OH⁻ < H₂O < NH₃ < en < CN⁻ < CO
- Weak field ligands (e.g. Cl⁻) cause small splitting → absorb lower energy (longer wavelength) light.
- Strong field ligands (e.g. CN⁻, NH₃) cause large splitting → absorb higher energy (shorter wavelength) light.
Example:
- [Cr(H₂O)₆]³⁺ is green/violet (water ligands).
- [Cr(NH₃)₆]³⁺ is purple (ammonia ligands cause greater splitting, shifting the absorbed wavelength).
2. The Oxidation State of the Metal
A higher oxidation state means:
- Greater charge on the metal ion.
- Stronger attraction for ligand lone pairs.
- Ligands drawn closer → greater splitting Δ.
- Different wavelength absorbed → different colour.
Example:
- [Fe(H₂O)₆]²⁺ is green (Fe is +2).
- [Fe(H₂O)₆]³⁺ is yellow/brown (Fe is +3; greater charge causes greater splitting).
3. The Coordination Number / Geometry
Changing the number and arrangement of ligands changes the splitting:
- Octahedral complexes have larger splitting than tetrahedral complexes.
- This means octahedral and tetrahedral complexes of the same metal and ligand absorb different wavelengths.
Example:
- [Co(H₂O)₆]²⁺ is pink (octahedral, coordination number 6).
- [CoCl₄]²⁻ is blue (tetrahedral, coordination number 4).
The difference in both the ligand and the geometry causes a dramatically different colour.
Why Some Compounds Are Colourless
A compound will be colourless if d–d transitions cannot occur: