Coloured compounds

We observe varying frequencies of electromagnetic radiation in the visible region of the electromagnetic spectrum as colour. Different colours result from the changed energy distribution of photons (light quanta) after it has been reflected, transmitted or absorbed after hitting a substance. Because of their structure, transition metals form many different coloured ions and complexes. Colour even varies between the different ions of a single element - MnO₄⁻ (Mn in oxidation state 7+) is a purple compound, whereas Mn²⁺ is pale-pink.

Coordination by ligands can play a part in determining color in a transition compound, due to changes in energy of the d orbitals. Ligands remove degeneracy of the orbitals and split them into higher and lower energy groups. The energy gap between the lower and higher energy orbitals will determine the color of light that is absorbed, as electromagnetic radiation is only absorbed if it has energy corresponding to that gap. When a ligated ion absorbs light, some of the electrons are promoted to a higher energy orbital. When the electron returns back to the lower energy level, it can emit a photon of wavelength that is complementary to the wavelength absorbed. If this is all it does, there is no net absorption of wavelength and no color is observed (since the amount emitted cancels the amount absorbed). However, if an electron loses this absorbed energy as extra vibration (because of ligand interaction or by some other way), it now returns to the lower energy orbital having yielded a net absorption of wavelength. Now the color observed is the color complementary to that which was absorbed since what is observed is white light minus the wavelength absorbed. For example, if blue is color observed, that implies that the wavelength of orange length was absorbed as an electron was excited up to the higher energy level, because orange is the complementary color to blue in white light. This energy is then lost by stronger than usual vibration, thus yielding a net absorption and causing the electron to fall back to the lower energy orbital.

The colour of a complex depends on:

• the nature of the metal ion, specifically the number of electrons in the d orbitals
• the arrangement of the ligands around the metal ion (for example geometric isomers can display different colours)
• the nature of the ligands surrounding the metal ion. The stronger the ligands then the greater the energy difference between the split high and low 3d groups.

The complex ion formed by the d-block element zinc (though not strictly a transition element) is colorless, because the 3d orbitals are full - no electrons are able to move up to the higher group.