Why do transition metals make colored compounds despite the remaining d orbitals being hybridized? [closed]








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I know that transition metals make colored compounds when they forms a bond with a ligand, as the ligand increases the energy level of $\ce{d\_{x^2-y^2}}$ and $\ce{d\_{z^2}}$ orbital by repelling these orbitals with the lone pair electrons of the ligand. So, the electrons move to the lower energy orbitals and when they absorb energy, they again move to the higher orbitals and that absorbed energy creates color.


Now my question is, $\ce{d\_{x^2-y^2}}$ and $\ce{d\_{z^2}}$ orbital become hybridized while forming a bond with the ligands and those orbitals get filled up with the lone pair electrons of ligands. Then, how are the electrons excited from lower energy orbitals into the higher energy orbital where there is no place left for them?


