Colour in transition metal complexesEdexcel A-Level Chemistry: Revision notes

Most transition metal ions in solution are coloured, e.g. $[\mathrm{Cu(H_2O)_6}]^{2+}$ is pale blue, $[\mathrm{Co(H_2O)_6}]^{2+}$ is pink and $[\mathrm{Fe(H_2O)_6}]^{2+}$ is pale green. The colour depends on the metal, its oxidation number, the ligands and the coordination number.

Colour arises because the ion absorbs some wavelengths of visible light. The light that is transmitted (or reflected) is seen as the complementary colour of the light absorbed.

In an isolated ion the five 3d orbitals have the same energy. When ligands approach, their lone pairs repel electrons in the d orbitals, and the orbitals split into two groups of different energy (in an octahedral complex, three lower and two higher). The energy gap between them is $\Delta E$.

A d electron in the lower group can absorb a photon of visible light with energy exactly equal to $\Delta E$ and be promoted to the higher group (a d–d transition):

$\Delta E = h\nu = \frac{hc}{\lambda}$

Worked example. Absorption at 650 nm: $\Delta E = \dfrac{6.63\times10^{-34}\times 3.00\times10^{8}}{650\times10^{-9}} = 3.06\times10^{-19}$ J per ion.

White light contains all visible wavelengths. The ion absorbs the frequency matching $\Delta E$. The remaining wavelengths combine to give the colour you see, the complementary colour.

Example: $[\mathrm{Ti(H_2O)_6}]^{3+}$ absorbs green light (about 500 nm) and looks purple.

A larger $\Delta E$ means light of higher frequency (shorter wavelength) is absorbed. Anything that changes $\Delta E$ changes the colour.

An ion is colourless when no d–d transition is possible in the visible range.

• $\mathrm{Sc^{3+}}$ is $[\mathrm{Ar}]$ with an empty 3d subshell: no d electrons to promote.
• $\mathrm{Zn^{2+}}$ ($3d^{10}$) and $\mathrm{Cu^{+}}$ ($3d^{10}$) have a full 3d subshell: nowhere for an electron to be promoted to within the d subshell.

Such ions absorb no visible light, so all wavelengths are transmitted and the solution is colourless (or white as a solid).

A change in any of these changes $\Delta E$ and so the colour:

• Oxidation number: $[\mathrm{Fe(H_2O)_6}]^{2+}$ is pale green; Fe³⁺(aq) is yellow-brown.
• Ligand: $[\mathrm{Cu(H_2O)_6}]^{2+}$ is pale blue, but $[\mathrm{Cu(NH_3)_4(H_2O)_2}]^{2+}$ is deep blue.
• Coordination number (with a change of ligand and shape): $[\mathrm{Co(H_2O)_6}]^{2+}$ is pink and octahedral; $[\mathrm{CoCl_4}]^{2-}$ is blue and tetrahedral.

In each case the explanation is the same: a different $\Delta E$, so light of a different frequency is absorbed and a different colour is transmitted.