O)] Blue Green Red [Co(NH ) ] Blue Yellow Orange [Co(CN) ] Ultraviolet Pale Yellow [Cu(H O) ] Red Blue [Ti(H O) ] Blue Green Violet The colour in the coordination compounds can be readily explained in terms of the crystal field theory. Consider, for example, the complex [Ti(H O) ] + , which is violet in colour. This is an octahedral complex where the single electron (Ti + is a d system) in the metal d orbital is in the t 2g level in the ground state of the complex. The next higher state available for the electron is the empty e g level.
If light corresponding to the energy of blue-green region is absorbed by the complex, it would excite the electron from t 2g level to the e g level ( t 2g e g ® t 2g e g ). Consequently, the complex appears violet in colour (Fig. . ).
The crystal field theory attributes the colour of the coordination compounds to d-d transition of the electron. ( b ) Crystal field splitting in tetrahedral coordination entities In tetrahedral coordination entity formation, the d orbital splitting (Fig. . ) is inverted and is smaller as compared to the octahedral field splitting.
For the same metal, the same ligands and metal-ligand distances, it can be shown that D t = ( / ) D . Consequently, the orbital splitting energies are not sufficiently large for forcing pairing and, therefore, low spin configurations are rarely observed. The ‘g’ subscript is used for the octahedral and square planar complexes which have centre of symmetry. Since tetrahedral complexes lack symmetry, ‘g’ subscript is not used with energy levels.
Not in visible region It is important to note that in the absence of ligand, crystal field splitting does not occur and hence the substance is colourless. For example, removal of water from [Ti(H O) ]Cl on heating renders it colourless. Similarly, anhydrous CuSO is white, but CuSO .5H