Undergraduate · 20 min

Why copper(II) complexes change colour

Switch between $\ce{[Cu(H2O)6]^{2+}}$ , $\ce{[Cu(NH3)4]^{2+}}$ and $\ce{[CuCl4]^{2-}}$ and link each colour to the ligand-field splitting $\Delta$ .

Goal

Rank the ligands $\ce{Cl^-}$ , $\ce{H2O}$ , $\ce{NH3}$ by field strength using the colour each complex shows.

Apparatus and reagents

Three flasks of copper(II) solution in different ligand environments, plus the spectrochemical series as reference.

Procedure

Select the water mode: the hexaaqua ion is pale blue. Note which wavelength it absorbs.
Switch to ammonia: excess $\ce{NH3}$ replaces water ligands and the colour becomes deep blue-violet.
Switch to concentrated $\ce{HCl}$ : the tetrahedral $\ce{[CuCl4]^{2-}}$ ion is green-yellow. Compare all three colours.

What to observe

$\ce{[Cu(H2O)6]^{2+}}$ pale blue → $\ce{[Cu(NH3)4]^{2+}}$ deep violet-blue → $\ce{[CuCl4]^{2-}}$ green-yellow.
The absorbed wavelength shifts from ≈ 800 nm (weak field) toward ≈ 600 nm (stronger field), i.e. $\Delta$ grows along $\ce{Cl^-} < \ce{H2O} < \ce{NH3}$ .

Explanation

A $d^9$ Cu(II) ion in a ligand field absorbs one photon whose energy equals the d–d splitting: $E = h\nu = \Delta$ . Strong-field ligands give a large $\Delta$ , so absorption shifts to shorter wavelengths and the observed (complementary) colour changes: pale blue for $\ce{[Cu(H2O)6]^{2+}}$ ( $\lambda_{abs} \approx 800$ nm), deep blue-violet for $\ce{[Cu(NH3)4]^{2+}}$ ( $\approx 600$ nm), green-yellow for tetrahedral $\ce{[CuCl4]^{2-}}$ whose splitting is only $\Delta_t \approx \frac{4}{9}\Delta_o$ . This ordering is the spectrochemical series.

Chemists behind it

Alfred Werner