One of the remarkable properties of transition elements is their color. It has been seen that most of the transition metal compounds show particular colors. This means that some of the visible spectra are absorbed by these elements from white light as it passes through a sample of transition metals. When transition elements are not bonded to anything else, their d orbitals are degenerate, that is, they all have the same energy level.
Why do we see different colors among transition elements?
When they start bonding with other ligands, due to different symmetries of the d orbitals and the inductive effects of the ligands on the electrons, the d orbitals split apart and become non-degenerate. When an electron jumps from lower energy d orbital to higher energy d orbital, that is a d-d transition, the energy of excitation corresponds to the frequency of light absorbed.
Thus, the energy required by the electrons for a change is provided by the light waves.The frequency of a light wave is observed to lie invisible range. The frequency of light absorbed depends on the nature of ligands. For example, if the electrons in an octahedral metal complex can absorb green light and get promoted from the dyz orbital to the dz2 orbital, the compound will reflect all the colors except green. Hence, the complementary color of green will be observed as the color of the compound.
From the above picture, we can easily visualize the different energy levels of the d orbitals. Therefore, an excitement of an electron from lower energy level to higher energy level requires energy. Hence, we can also conclude that not all transition metal complexes are colored as transition elements with fully filled d orbital do not allow the possibility of d-d transitions. Hence, no radiations are absorbed. For example, Zinc Sulphate.
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