IBDP Chemistry: AHL Topic 13.2-Coloured complexes: Study Notes

13.2 Coloured complexes

Essential Idea:
d-orbitals have the same energy in an isolated atom, but split into two sub-levels in a complex ion. The electric field of ligands may cause the d-orbitals in complex ions to split so that the energy of an electron transition between them corresponds to a photon of visible light.

Understandings:

  • The d sub-level splits into two sets of orbitals of different energy in a complex ion.
  • Complexes of d-block elements are coloured, as light is absorbed when an electron is excited between the d-orbitals.
  • The colour absorbed is complementary to the colour observed.

Applications and Skills:

  • Explanation of the effect of the identity of the metal ion, the oxidation number of the metal and the identity of the ligand on the colour of transition metal ion complexes.
  • Explanation of the effect of different ligands on the splitting of the d-orbitals in transition metal complexes and colour observed using the spectrochemical series.

13.2 Coloured complexes – Facts

Transition metal ions are coloured due to d–d electron transitions between d orbitals which are split in the electric field due to the presence of the ligands.

• The colour observed is complementary to the colour absorbed, and can be deduced from the colour wheel.
• The colour of a complex depends on the identity of the metal ion, the oxidation state of the metal, and the identity of the ligand.
• Ions with higher charge and ligands with greater charge density produce a greater split in the d orbitals.
• The spectrochemical series arranges the ligands according to the energy separation between the two sets of d orbitals.
• Polydentate ligands form more than one coordinate bond with the metal ion.

d sub-level splitting

A compound that contains energy levels that are close together could absorb radiation in the visible light spectrum and therefore display a colour because the colour observed is complementary to the colour absorbed.  This must be the case in d-block metals as their compounds are frequently coloured both in solid state (hydrated) and in solution.

In the d-block metal ions in complexes in their ground state, the five orbitals in the 3d sub-level have the same energy (equivalent energy level) but in most of their compounds (usually complexes), these 5 equivalent orbitals are split in two or three different energy level (non-equivalent energy levels) with some orbitals on each level.

This splitting of the 3-d sublevel into two or more sets of orbitals is caused by the ligands as the electron clouds or non-bonding pairs around the ligand repel the electrons in the 3d orbitals of the metal ion.  As a result the ligand electrons push the 3d electrons in the transition metal ion in orbitals closest to them to energy levels higher than the other orbitals that are not near the ligand.  This creates sets of orbitals of different energy – non-equivalent – and this process is called field-splitting.

d-to-d transition

Although they are different in energy, the two sets of split orbitals are still close together.  This allows an electron in a lower 3d orbital to become excited and absorb radiation (from the visible light spectrum) and be promoted (=transition) from a lower 3d orbital (low spin) to a higher 3d orbital (high spin).  The amount of energy needed for this d-to-d transition (or the energy difference between the two sets of orbitals) corresponds to a photon within the visible light spectrum. 

The colour shown by the transition metal ion complex is a mixture of the colours or radiation that it transmits after absorption of the frequencies for transition. The colour observed is a mixture of the colours complementary to those colours that have been absorbed.  You should use a colour wheel to determine the complementary colour of a colour that has been absorbed.

Therefore for transition metal compounds to form coloured compounds, the transition metals ions must have partially-filled 3 d orbitals i.e. unpaired electrons.

The field splitting in a Cu2+ when it forms a complex with water is shown in the diagram below (from http://www.chemguide.co.uk/inorganic/complexions/colour.html). The symbol ΔE indicates the difference in energy between the two sets of orbitals. Using Planck’s constant and the frequencies of the complementary colours you should be able to calculate the amount of energy.

The identity of the ligand: each ligand has its own effect on the relative energies of the d electrons, e.g. NH3 has a greater effect than water because nitrogen has a lower electronegativity and therefore attracts its lone pair less strongly/less closely to the nucleus allowing it to repel more other electrons.  NH3 is considered a stronger ligand than water. The stronger the ligand the more the electrons in the 3d split sub-level absorb towards the high-energy end of the visible light spectrum. Electronegativity cannot always be used to explain the difference as Cl is a stronger ligand than I as it has a higher charge density than I (same ionic charge but greater ionic radius).

Spectrochemical series

The spectrochemical series is a series that ranks various ligands in order of increasing ΔE values (large splitting). The series is based on experimental evidence. The further up the series the ligand, the more splitting of it causes, the higher the frequency/lower the wavelength of the radiation needed for a d-to-d transition.  

Example of the effect of a ligand on the colour observed: hexaaquacopper (II) ions have a blue colour. As ammonia is a stronger ligand it displaces four of the water molecules and converts the complex ions into deep blue tetraaminediaquacopper (II) ions. This happens because ammonia as a stronger ligand causes greater repulsion and therefore a greater ΔE.  As a result electrons in the tetraaminediaquacopper (II) complex need to absorb light of a higher energy/frequency to make any d-to-d transitions than in the hexaaquacopper (II) complex. This means the complementary colours that are observed also have higher frequencies.

Not all d-block compounds are coloured compounds

Compounds containing transition ions with empty d-orbitals (e.g. Sc3+, Ti4+) or full d orbitals (e.g. Zn2+) are colourless as no transition between split sub-levels or different sets of orbitals can occur (in the case of full orbitals there are no spaces). 

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