Colour of Transition metals

  • Colours and colour changes are among the most striking aspects of the d-block transition metals.
  • Both in solid state and in solution, the transition metals show a great variety of colours.
  • The colours are formed due to the presence of partially filled d-orbitals.
  • When light falls on an object, some of it is absorbed, some transmitted and the rest reflected; this results in colours being produced.
  • Light is only absorbed when its energy matches the difference between two energy states in the atom.
  • If it does, an electron is promoted from a lower energy orbital to a higher one (i.e. atom is excited from ground state to an excited state).
  • For an uncombined metal ion (gaseous), the five d-orbitals have the same energy.
  • But in a complex, the metal ion is surrounded by ligands with lone pairs of electrons.
  • The presence of the ligands affects the electrons in the d-orbitals; the electrons in the orbitals closer to the ligands are pushed to a higher energy level; this results in the five d-orbitals being split into two energy levels.
  • The energy difference between these two groups of energy levels (in most transition metals) is equivalent to a frequency of light within the visible region of the electromagnetic spectrum (related to E = h); therefore the complex appears to be coloured.


The graph below shows the relative energy levels for the 5 d-orbitals of hydrated Ti3+ ion:

  • E corresponds to the green/yellow frequency of light; the green/yellow light is absorbed by the atoms. This results in the Ti 3+ solution appearing purple in colour.
  • The example above is only simple; there is only one electron in the d orbital of Ti.
  • When there are several d-orbital electrons, the situation is more complex as there can be many transitions of electrons between the energy levels.
  • The colour of a transition metal complex depends upon:
    • The number of d-electrons present.
    • The arrangement of the ligands around the ion (different arrangements affect the splitting of the orbital).
    • The nature of the ligand; different ligands have different effects upon the relative energies of the d-orbitals. For example, NH3 ligands cause a greater difference than H2O ligands in splitting d-orbital energies; thus the colour changes from green/blue to deep blue/violet when NH3 is added to aqueous Cu(ii) salt.
    • The number of ligands and the way that they are arranged around the ion also has an effect.

    Useful books for revision

    Revise A2 Chemistry for Salters (OCR A Level Chemistry B)
    Salters (OCR) Revise A2 Chemistry