5.5.5. Colour in Coordination Compounds

In tetrahedral coordination entity formation, the d orbital splitting 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 Dt = (4/9) D0. 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.

In coordination compounds, the arrangement of ligands around a metal center affects the energy levels of its d orbitals. In octahedral complexes, d orbitals split into two energy levels due to the symmetrical placement of ligands. Tetrahedral complexes have a different splitting pattern where the lower energy orbitals are now higher in energy compared to the octahedral case. This results in a smaller energy difference (Dt) in tetrahedral complexes, making it less likely for electrons to pair up, often leading to high spin states.

Think of a room filled with chairs (ligands) set up for a meeting. In an octahedral arrangement, chairs are placed symmetrically around a central table (the metal). This setup naturally divides the space into 'boring' and 'exciting' areas. In a tetrahedral setup, the chairs are arranged unevenly, making the 'boring' space feel less appealing, leading to people (electrons) staying rather than pairing up with someone else (which would cost energy).

In the previous Unit, we learnt that one of the most distinctive properties of transition metal complexes is their wide range of colours. This means that some of the visible spectrum is being removed from white light as it passes through the sample, so the light that emerges is no longer white. The colour of the complex is complementary to that which is absorbed. The complementary colour is the colour generated from the wavelength left over; if green light is absorbed by the complex, it appears red. Table 5.3 gives the relationship of the different wavelength absorbed and the colour observed.

Transition metal complexes can absorb specific wavelengths of light due to the excitation of electrons between different d orbitals. The remaining wavelengths that are not absorbed are what we perceive as the color of the compound. For instance, if a compound absorbs light in the green range of the visible spectrum, it will appear red, which is the complementary color. This interaction is described by the color wheel, where opposite colors enhance our understanding of how these compounds behave.

Imagine wearing sunglasses that only let red light pass through. You would see everything around you in shades associated with that filter's effect, much like how transition metal dyes absorb some colors of light and reflect others. If you put on green-tinted glasses that filter out reds, you'll see all other colors shifted, which is akin to how color in coordination compounds works when certain wavelengths are absorbed.

The colour in the coordination compounds can be readily explained in terms of the crystal field theory. Consider, for example, the complex [Ti(H2O)6]3+, which is violet in colour. This is an octahedral complex where the single electron (Ti is a 3d system) in the metal d orbital is in the t2g level in the ground state of the complex. The next higher state available for the electron is the empty eg level. If light corresponding to the energy of blue-green region is absorbed by the complex, it would excite the electron from t2g level to the eg level (t2g eg ® t2g eg). Consequently, the complex appears violet in colour.

In this specific complex, the presence of water ligands creates an octahedral arrangement around the titanium ion. The electron transitions occurring within the split d orbitals cause the complex to absorb specific wavelengths of light, primarily those associated with blue-green. The energy absorbed causes an electron to jump to a higher energy level, changing the way we perceive the color of the complex, making it appear violet.

Consider a prism separating white light into a rainbow. Each colored band corresponds to different wavelengths of light. When a specific band of light is absorbed by a chemical substance, the remainder determines how we see it. Just as different prisms can filter out specific colors, in compounds like [Ti(H2O)6]3+, certain light wavelengths get absorbed for electron transitions, letting us see the color that remains, much like the colors of a sunset filtered through clouds.

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