Exciting electrons from lower to higher energy d orbitals corresponds to the absorption of light, with the absorbed light’s frequency determining the observed color. The nature of the ligand, particularly in aqueous solutions where water molecules act as ligands, influences the frequency of absorbed light.
How does the absorption of light, corresponding to the excitation of electrons from lower to higher energy d orbitals, result in the observation of colors, and what determines the frequency of light absorbed in the visible region?
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The absorption of light, corresponding to the excitation of electrons from lower to higher energy d orbitals in transition metal complexes, leads to the observation of colors. This excitation involves the promotion of electrons from the ground state to an excited state. The frequency of light absorbed in the visible region is determined by the energy difference between the lower and higher energy orbitals. The color perceived is complementary to the absorbed wavelength. Complexes absorbing longer wavelengths appear red, while those absorbing shorter wavelengths appear violet. This relationship between electron transitions and absorbed light frequencies defines the visible color spectrum.