maandag 13 oktober 2025
zondag 5 oktober 2025
The Colourful Coordination Complexes of Copper II – Absorption Spectrometry
Copper is a d-group metal with a full 3d orbital: the electron configuration of the neutral atom is 4s1 3d10. When copper is oxidised to Cu2+ two electrons are absorbed by the oxidiser, leaving us with the electron configuration 3d9. The near-full 3d configuration opens up the possibility of coloured coordination complexes.
The full 3d orbital is made up of five suborbitals, each of which can ‚house‘ 2 paired electrons. It can thus accomodate 10 (= 5 x 2) electrons.
The shape of the five d-orbitals is shown below (schematically):
(a much better representation, based on actual wave function calculations can be found at the Orbitron)
The d-(sub)orbitals can be classed into two groups:
* the e-group, which have their strongest electron densities along the x, y and z axis – these are the 2 orbitals in the top row
* the t-group, which have their strongest electron densities along the xy, yz and xz planes – these are the 3 orbitals in the bottom row
It is important to note that the five d orbitals are so-called ‚degenerate‘, meaning that their energy levels are identical.
Now, enter one or more ligands (say L), attracted by the central electrical field of the Cu2+ ion, along the x,y and z principal axis. The ligands have their own electron densities and these perturb the electron densities of the suborbitals, through electrostatic repulsion. However, the e-group suborbitals undergo more perturbation by the approaching ligands L, than the t-group orbitals.
This creates a difference in energy levels between the e and t orbitals, which I will call ΔE.
The new ground state of the system now becomes (left no ligands - right ligands present):
This now opens the possibility to excite one of the lower energy t-electrons and kick it into one of the half empty e-suborbitals. The energy ΔE needed for this is, depending on the Ligand L, in the order of about 2 eV (electron volt), which happens to lie in the visible range of the electromagnetic spectrum.
So photons in the 400 nm (violet) to 800 nm (red) range can kick a t-electron into an e-orbital. The photon is thus absorbed by the coordination complex.
This is of course what we observe: copper (and many other d-block metals) form richly coloured coordination complexes, with a variety of ligands, because they absorb photons in the visible spectrum. Here is what a few of them look like for different ligands:
A little Absorption Spectrometry
Firstly, the reference spectrum, generated with a powerful halogen lamp. This is used to irradiate the complexes:
1. hexaqua complex ([Cu(H2O)6]2+):
There's a striking, almost 100 % absorbance, from about the orange all the way up to the infra-red (the black tail on the right hand side)
2. tetra-ammonia complex ([Cu(NH3)4]2+):
Not dissimilar to the aqua complex (perhaps because both ligands are electrically neutral)
3. tetrahydroxy complex ([Cu(OH)4]2-):
This complex is a bit dissapointing, probably because of its low concentration. I may have to repeat it.
4. tetrachloro complex ([CuCl4]2-):
Interestingly this complex appears almost 100 % transparent for green light only.
