This graphic looks at the colours of transition metal ions when they are in aqueous solution (in water), and also looks at the reason why we see coloured compounds and complexes for transition metals. This helps explain, for example, why rust (iron oxide) is an orange colour, and why the Statue of Liberty, made of copper, is no longer the shiny, metallic orange of copper, but a pale green colour given by the compound copper carbonate.
In order to explain why transition metals are coloured, we first have to talk a little about how the electrons in an atom are arranged around the central nucleus. In secondary school, the majority of students learn that electrons are arranged in ‘shells’ around the nucleus; whilst this is a useful model for looking at electron arrangements, there is also an extra layer of complexity.
Within shells, electrons are actually arranged in special areas in particular energy levels, in sub-shells called ‘orbitals’. These orbitals come in different shapes, and are named using different letters: s, p, d, & f. Each of this orbitals can hold varying numbers of electrons: s can hold 2, p 6, d 10 and f 14. Transition metals are unique in the Periodic Table in that they are the only elements that contain partially filled d orbitals, and these are key to the coloured compounds and complexes they form.
Transition metal complexes are formed when transition metals are bonded to one or more neutral or negatively charged non-metal species, referred to as ‘ligands’. Without these bonds, all the d orbitals are equal in energy – however, once they are present, some d orbitals move to a higher energy than they were at before, whilst some move to a lower energy, creating an energy gap. This is due to the fact that, due to their differing shapes, some d orbital are nearer to the ligands than others. Electrons can move from the lower energy d orbitals to the higher energy d orbitals by absorbing a photon of light; the wavelength of the absorbed light depends on the size of the energy gap. Any unabsorbed wavelengths of light pass through unabsorbed, and this causes the coloured appearance of the compounds.
The colour can be affected by several variables. Different transition metals will exhibit different colours; as shown in the graphic above, different charges on the same transition metal can also accomplish this. The ligand also has an effect, and the same charge metal ion can be differently coloured depending on the ligands that are bound to it.
Alternatively, you can download this graphic in A3 PDF format here.
References & Further Reading