Coloured Complexes (Cambridge (CIE) A Level Chemistry)

Coloured Compounds & Electron Promotion

• Most transition element complexes are coloured

• A transition element complex solution which is coloured, absorbs part of the electromagnetic spectrum in the visible light region

• The observed colour is the complementary colour which is made up of light with frequencies that are not absorbed

  • For example, copper(II) ions absorb light from the red end of the spectrum

  • The complementary colour observed is therefore pale blue (cyan)

The visible light region of the electromagnetic spectrum

The visible light region ranges from red to violet

Electron promotion

• In an isolated transition element ion (which is not bonded to any ligands), all of the 3d orbitals are degenerate

• However, when ligands are attached to the central metal ion through dative covalent bonds, these orbitals are split into two sets of non-degenerate orbitals

• The difference in energy between these two sets of orbitals is ΔE

• When light shines on a solution containing a transition element complex, an electron will absorb this exact amount of energy (ΔE)

• The amount of energy absorbed can be worked out by the equation:

ΔE = h x v

• Where:

  • h = Planck's constant (6.626 x 10^-34 m² kg s^-1)

  • v = frequency (Hertz, Hz or s^-1)

• The electron uses the energy from the light to jump into a higher, non-degenerate energy level

  • This is also called electron promotion

• The other frequencies of light which are not absorbed combine to make the complementary colour

• The diagram below shows an example of electron promotion in an octahedral complex of a nickel(II) Ni²⁺ ion

Electron promotion in a Ni(II) complex when light shines on the solution

An electron gains enough energy to be promoted from a lower energy non-degenerate orbital to a higher energy non-degenerate orbital

Effects of Ligands on Complementary Colour

• Transition element complexes absorb the frequency of light which corresponds to the exact energy difference (ΔE) between their non-degenerate d orbitals

• The frequencies of light which are not absorbed combine to make the complementary colour of the complex

• It is the complementary colour which is seen

• However, the exact energy difference (ΔE) is affected by the different ligands which surround the transition element ion

• Different ligands will split the d orbital by a different amount of energy

• This depends on the repulsion that the d orbital experiences from these ligands

• Therefore, the size of ΔE and thus the frequency of light absorbed by the electrons will be slightly different

• As a result, a different colour of light is absorbed by the complex solution and a different complementary colour is observed

• This means that complexes with similar transition elements ions, but different ligands, can have different colours

  • For example, in copper complexes:

  • [Cu(H₂O)₆]²⁺ complex has a light blue colour

  • [Cu(NH₃)₄ (H₂O)₂]²⁺ has a dark blue colour

  • Despite the copper ion having an oxidation state of +2 in both complexes

  • This is evidence that the ligands surrounding the complex ion affect the colour of the complex

Ligand Exchange in Copper(II) & Cobalt(II) Complexes

• Different ligands may affect the complementary colour of a transition ion complex solution

• This is shown by ligand exchange reactions in copper(II) and cobalt(II) complexes, as this causes a change in colour of the complexes

Copper(II) & cobalt(II) ions

• The ligand exchange of [Cu(H₂O)₆]²⁺ and [Co(H₂O)₆]²⁺ by NH₃ ligands causes a change in the colour of the solutions

  • [Cu(H₂O)₆]²⁺ is light blue in colour whereas [Cu(NH₃)₄(H₂O)₂)]²⁺ is deep blue in colour

  • [Co(H₂O)₆]²⁺ is a pink solution whereas [Co(NH₃)₆]²⁺ is a brown solution

• The colour change results from the ammonia ligands, which cause the d orbitals to split by a different amount of energy (ΔE)

• Therefore, the size of ΔE and the frequency of light absorbed by the electrons will be slightly different

• As a result, a different colour of light is absorbed and thus a different complementary colour is observed

Ligand exchange of hexaaqua copper(II) by ammonia

Ligand exchange of the water ligands by ammonia ligands causes a change in the colour of the copper(II) complex solution from blue to dark blue

Ligand exchange of hexaaqua cobalt(II) by ammonia

Ligand exchange of the water ligands by ammonia ligand causes a change in the colour of the cobalt(II) complex solution from pink to brown

• Similarly, full ligand exchange by chloride ions in copper(II) and cobalt(II) complexes results in a change in complementary colour

  • [Cu(H₂O)₄(OH)₂] is a pale blue precipitate whereas [CuCl₄)]^2– is a yellow solution

  • [Co(H₂O)₄(OH)₂] is a blue precipitate whereas [CoCl₄)]^2– is a blue solution

• The colour change results from the chloride ligands, which cause the d orbitals to split by a different amount of energy (ΔE)

• Therefore, the size of ΔE and the frequency of light absorbed by the electrons will be slightly different

• As a result, a different colour of light is absorbed and thus a different complementary colour is observed

Ligand exchange of [Cu(H₂O)₄(OH)₂] by chloride ions

Ligand exchange by chloride ligands causes a colour and state change in the colour of the copper(II) complex from a pale blue precipitate to a yellow solution

Ligand exchange of [Co(H₂O)₄(OH)₂] by chloride ions

Ligand exchange by chloride ligands causes a colour and state change in the colour of the cobalt(II) complex from a blue precipitate to a blue solution

• As before, this suggests that different ligands will split the d orbitals differently