[hale2bopp]

When a strong field ligand forms a complex with a metal, the energy gap between the eg and t2g levels is quite high. So, once the complex has been formed, if we make light incident on the complex, it would absorb light of high frequency since the energy gap is large. Since frequency of light absorbed is high, wavelength is low. Correspondingly, the wavelength of the colours of light left over would be high and this is the colour that reaches us. Since the absorbed wavelengths are low, the leftover wavelengths would be high, meaning that the colour reaching us would be reddish. Then, how come, when we have a nickel complex, and we add ethane-(1,2)-diamine, which is a strong field ligand, it turns blue and then purple?

[blaisem]

en is a strong ligand, but nickel forms a square planar complex, which has only two degenerate d-orbitals.

A d⁸ metal, its transitions aren't as straightforward as promoting an electron from the ground state and moving to the highest energy orbital, as in octahedral.  If you look at a picture of the square planar d orbital energies, an electron can absorb from the d_xy to the highest energy level - d_x²-y², which is not the total energy of delta - it is less and therefore it absorbs apparently a yellow green light.

Picture from Wikipedia crystal field theory page.

[cheese (MSW)]

F. A. Cotton, G. Wilkinson, C. A. Murillo, M. Bochmann,
Advanced Inorganic Chemistry 6th ed (1999). p 838-839
Explanation of why [Ni(en)3]^2+ (i.e., octahedral) is blue-violet.