Colour of Coordination Metal Complexes | Inorganic Chemistry PDF Download

COLOUR

A characteristic feature of many d-block metal complexes is their colours, which arise because they absorb light in the visible region. For example, [Cr(H₂O₆]²⁺ is sky blue, [Mn(H₂O)₆]²⁺ very pale pink, [Co(H₂O)₆]²⁺ pink, [MnO₄]^- intense purple and [CoCl₄]^2- dark blue. In contrast, salts of Sc(III) (d⁰) or Zn(II) (d¹⁰) are colourless.

Colour in d-block complexes arise from transitions between electronic energy levels:

• Transitions between metal-centred orbitals possessing d character (‘d–d’ transitions). This transition takes place between t_2g and eg orbitals.
•  Transitions between metal and ligand centred MO’s which transfer charge from metal to ligand or ligand to metal, also called charge transfer spectra.

Charge transfer (CT) gives rise to intense absorptions, whereas ‘d–d’ bands are much weaker.
Using the following colour wheel we can determine the observed colour of a compound from the colour of absorbed light.

In the following colour wheel, the complementary colours are given in opposite sectors to each others. For example, if any complex is absorbing red colour light, it will transmit green colour and hence will look green. If it will absorb violet colour, it will be yellow in colour and so on.  

In some complexes the ∆ value is very small and these absorb light in infrared region and hence are colourless. Similarly, the complexes which have large ∆ value absorb radiations in ultraviolet region and hence appear colourless. The intensities of all absorption bands are not same. For example, when d-d transition occurs in complexes with centre of symmetry such as regular Oh complexes, the intensities of absorption bands are low. But when d-d transition occurs in complexes which lack centre of symmetry such as tetrahedral complexes and Oh cis complexes of type MA₄B₂, relatively strong absorption bands are observed.

Selection Rules: The variation in relative intensities of absorption bands are explained by selection rules:

• Laporte’s Selection Rule: The transitions which led to change in azimuthal quantum no. are allowed transitions i.e. transitions are Laporte allowed if,
  ∆l = ±1

Where l is the azimuthal quantum number of the orbitals between which transition takes place.
If ∆l = 0, the transitions are Laporte forbidden and therefore, absorbance is low. This rule can also be formulated in another way as following.
The transition that occur between states of opposite parity i.e. g  u are allowed. For example, s  p , p d and d  f transitions are allowed.
The transitions that occur between states of same parity i.e. between g  g or u  u are forbidden. For example, d-d or f-f transitions are forbidden because d and f-orbitals are gerade.

• Spin Selection Rule: Transitions between states of same multiplicit y are allowed. i.e.

If ∆S = 0, spin allowed.
In other words, during electronic transitions if the spin of the electron is not changed, then the transition is spin allowed.  Transitions between states of different multiplicities are spin forbidden i.e.

If ∆S ≠ 0, spin forbidden.
For example, Mn²⁺ complexes give ver y less intense colour because the transition in these complexes from t_2g to eg is spin forbidden.

A Laporte forbidden d-d transition is more intense than a spin forbidden transition. The intensities of various spectroscopic bands of 3d metal complexes are given below:

From the above table we see that charge transfer bands are most intense and spin forbidden transitions are least intense.

Relaxation in Selection Rules: The d–d transitions are forbidden, even some complexes gives colours. This is due to the relaxation in above selection rules. For example, 

•  The metal-ligands bonds in transition metal complexes are not rigid. When UV light is incident on a complex, electronic transitions as well as vibrations occur simultaneously. These vibrations destroyed centre of symmetry temporarily and there is small mixing of p-d orbitals and thus d-d transitions are not purely Laporte forbidden. Therefore, octahedral complexes show colour of low intensity.
• In tetrahedral (Td) complexes there is no centre of symmetr y and p -d mixing is more pronounced in Td complexes. Therefore, Td complexes give more intense colour than Oh complexes.
• There is some relaxation of spin selection rules in complexes having spin-orbit coupling. The strength of spin orbit coupling in lighter atoms is lower than for heavier atoms.  For example, spin-orbit coupling is strong in 4d and 5d transition metal complexes than in 3d metal complexes. Therefore, these complexes give intense colour than 3d metal complexes.

The above statements is summarised in the table given below:

Charge Transfer Spectra

Charge transfer transition is the transfer of an electron between orbitals of different atoms. These transitions are identified by their high intensity and the sensitivity of their energies to solvent polarity. This is due to that these transitions are Laporte and spin allowed i.e.
In a CT transition, an electron migrates between orbitals that are predominantly ligand in character and orbitals that are predominantly metal in character. When these transitions occur in visible region, the compounds show intense colour. A charge transfer is regarded as an internal redox process. Charge transfer spectra are of four types:

1. Ligand to Metal Charge Transfer (LMCT)
2. Metal to Ligand Charge Transfer  (MLCT)
3. Metal to Metal Charge Transfer
4. Intra Ligand Charge Transfer

The CT character is most often identified (and distinguished from π to π* transitions on ligands) by demonstrating solvatochromism, the variation of the transition frequency with changes in solvent permittivit y. Solvatochromism indicates that there is a large shift in electron densit y as a result of the transition, which is more consistent with a metal-ligand transition t han a ligand-ligand or metal-metal transition.

The following figure summarizes the transitions we classify as charge transfer:

1. Ligand to Metal Charge Transfer (LMCT): If the transfer of an electron takes place from the ligand to metal, then the charge transfer is called ligand to metal charge transfer. The conditions for LMCT are:

• Metal should be in high oxidation state, smaller energy and at lower energies.
• Ligands should have lone pair of electrons of relatively high energy and low electron affinit y.

For example, tetraoxidoanions of metals with high oxidation numbers such as MnO₄^– are most familiar examples of LMCT. In these, an O lone-pair electron is promoted into a low-lying empty metal e orbital. High metal oxidation numbers correspond to a low d-orbital population (many are formally d⁰), so the acceptor level is available and low in energy.
Energy required to transfer an electron from ligand to metal depends upon the LUMO of the metal ion and the HOMO of the ligand. For example, the trend in LMCT energies is: 

Oxidation number 
+7                                      MnO₄^– < TcO₄^– < ReO₄^–

+6                                      CrO₄^– < MoO₄^– <WO₄^–

+5                                      VO₄^– < NbO₄^– < TaO₄^–

This is due to that on moving from 3d to 4d to 5d series of transition metals in a group, size of metal ion increases and energy of LUMO increases. Therefore, the transfer of electron from HOMO to LUMO is feasible.

On moving from VO₄^3– to CrO₄^2– to MnO₄^– size of metal cation decreases. Therefore, energy of transfer decreases.
 
  

2. Metal to Ligand Charge Transfer (MLCT): In MLCT, an electron migrates from metal to ligand. MLCT are favoured in complexes in which:

• Metal have low oxidation state.
• Metal d orbitals are field and of high in energy.
• Ligand have empty π*-antibonding orbitals for example aromatic ligands. That is why MLCT mainly occurs in ligands having π* orbitals such as CO, CN^-, SCN^-, pyridine, bipyridine, o^-phenanthroline, p yrazine, dithiolene, NO etc.

In an octahedral complex when t_2g and e_g orbitals are occupied, two MLCT bands t_2g to π* and e_g to π* are observed.

Examples of compounds which show MLCT are given below:

Compound                                    Colour 
K₄ [Fe(CN)₆]                                  Yellow
K₃ [Fe(CN)₆]                                   Red
[Fe(phen)]³⁺                                   Blue
[Fe(acac)₃]                                      Red

3. Metal to Metal Charge Transfer (Intervalance Charge Transfer): In these transitions electron transfer takes place from metal of lower oxidation state to metal of higher oxidation state in a complex. For example,

(a) Prussian blue KFe[Fe(CN)₆] shows intense blue colour because of transfer of an electron from Fe²⁺ to Fe³⁺. In Prussian blue, Fe²⁺ is octahedrally coordinated with C of CN^– ligands and Fe³⁺ is octahedrally coordinated with N of CN^–. The electron transfer takes place through bridging cyanide ligand.

(b) Charge transfer in Creutz-Taube complex ion, [(NH₃)₅Ru-Pyz-Ru(NH₃)₅]⁵⁺, where p yz is bridging pyrazine ligand. 

In this complex ion electron transfer occurs from Ru(II) to Ru (III) through pyrazine bridging ligand and give intense colour.

(c) Red lead (Pb₃O₄) contains Pb(II) and Pb(IV). Due to the electron transfer from Pb²⁺ to Pb⁴⁺ this complex gives intense red colour.

4. Intra Ligand Charge Transfer: Some ligands itself behave as chromospheres. There are four transitions  σ → σ*, π→ π*, n→ π*, n→σ* occur in a chromophore. When such ligand is coordinated with metal ion/atom, complex show colour.