Colour of Complexes

Specification Reference Inorganic Chemistry, Chemistry of transition elements 28.3

Quick Notes

In octahedral complexes, d orbitals split into 3 lower and 2 higher energy levels.
In tetrahedral complexes, the split is reversed: 2 lower, 3 higher.
ΔE = energy gap between split d orbitals.
Light is absorbed when electrons jump from lower to higher d orbitals.
Colour observed is the complementary colour of the light absorbed.
Different ligands cause different splitting (ΔE), changing absorbed frequency and colour.
Ligand exchange can cause a visible colour change (e.g., [Cu(H₂O)₆]²⁺ vs. [Cu(NH₃)₄(H₂O)₂]²⁺).

Full Notes

d Orbital Splitting in Complexes

All five d orbitals (d_xy, d_yz, d_zx, d_z², d_x²−y²) have the same energy in a free transition metal ion and are described as degenerate.

When a complex forms, the electrons in the ligands repel the electrons in the d orbitals of the metal ion. This causes the d-orbitals to split into two sets with different energies – they are now non-degenerate. The pattern of splitting depends on the geometry of the complex ion:

Octahedral Complexes (e.g., [Fe(H₂O)₆]²⁺):

CIE A-Level Chemistry octahedral splitting of d orbitals into higher and lower energy groups.

The d orbitals split into two higher energy (d_x²−y² and d_z²) and three lower energy (d_xy, d_yz, d_zx) groups.
Electrons absorb energy ΔE to jump to the higher orbitals.

Tetrahedral Complexes (e.g., [CoCl₄]²⁻):

CIE A-Level Chemistry tetrahedral splitting of d orbitals showing inverted arrangement compared to octahedral.

The splitting pattern is opposite to octahedral splitting: three higher and two lower energy orbital groups.
The energy gap ΔE is smaller than in octahedral complexes.

Why Are Transition Metal Complexes Coloured?

Electrons absorb energy (ΔE) from visible light to move from lower to higher d orbitals.

CIE A-Level Chemistry diagram showing light absorption causing electron excitation between split d orbitals.

The frequency (or wavelength) of light absorbed depends on ΔE.
The colour observed is the complementary colour to the light absorbed.

CIE A-Level Chemistry colour wheel showing complementary colours in absorption and appearance of complexes.

Examples:

[Cu(H₂O)₆]²⁺ absorbs orange-red wavelengths of light and appears blue.
[Co(H₂O)₆]²⁺ absorbs yellow light wavelengths of light and appears pink.

Ligand Effects on ΔE and Colour

Different ligands cause different splitting energies (ΔE) depending on how strongly they interact with the metal ion:

Stronger field ligands (e.g., CN⁻, NH₃) cause greater splitting
higher frequency light absorbed (shorter wavelength).
Weaker field ligands (e.g., H₂O, Cl⁻) cause smaller splitting
lower frequency light absorbed (longer wavelength).

This means ligand exchange changes ΔE and can cause the colour of the complex to change.

Examples of Ligand Exchange and Colour Change

CIE A-Level Chemistry diagram showing ligand exchange effects and colour changes in metal complexes.

These colour changes are due to the difference in ligand field strength, which alters ΔE and changes the frequency of light absorbed.

Summary

In octahedral complexes, d orbitals split into groups of 3 lower energy orbitals and 2 higher energy orbitals (in tetrahedral complexes, the pattern is reversed).
Electrons absorb visible light energy (ΔE) to move between split d orbitals.
The observed colour is complementary to the absorbed wavelength.
Different ligands produce different splitting energies (ΔE), changing the absorbed frequency and colour.
Ligand exchange can cause visible colour changes by altering ΔE.