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Why do d-block elements have coloured complexes?

In complexes the d orbitals are split into two distinct levels, they are not degenerate (of equal energy) as they are in free ions. 

The energy difference between the levels corresponds to a particular wavelength or frequency in the visible region of the spectrum.

When light hits the complex, the energy of the corresponding wavelength is absorbed and electrons are excited from the lower energy level of d orbital to the higher level. 

The colour that is visible is the complimentary colour to the wavelengths that have been absorbed (e.g Cu2+(aq) appears blue because the complementary wavelengths of light have been absorbed.)

The colour of the complex depends on the energy difference between the d orbitals which in turn depends on several factors: the nature of the transition metal, the oxidation state, the shape of the complex and the nature of the ligand. This explains why different complexes are different colours.

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