Colour properties of nickel complexes

According to Crystal Field theory (CFT), complexes with stronger ligands must absorb light with higher frequency hence would transmit corresponding complementary colour.


The colour of $(\ce{[Ni(H2O)2(en)2]^2+})$ blue - purple and colour of $(\ce{[Ni(H2O)6]^2+})$ is green, whereas the colour of $(\ce{[Ni(en)3]^2+})$ purple. (Where $\ce{en ->}$ ethylene diamine.) (According to NCERT chemistry class 12)


[![enter image description here](https://i.stack.imgur.com/GkfWY.jpg)](https://i.stack.imgur.com/GkfWY.jpg)


But the above example goes against the CFT statement. Any plausible explanation available?




OP's argument:



> 
> According to Crystal Field Theory(CFT), complexes with stronger ligands must absorb light with higher frequency hence would transmit corresponding complementary color.
> 
> 
> 


First of all, the different colors we see in these solutions are not emission colors. These solutions do not emit as AChem pointed out in his comment. As he correctly put,
the colors are due to the transmitted light minus the absorbed portion of the spectrum. On other words, The color you see in s solution of a transition metal compound is due to the wavelengths of the visible light that isn't absorbed by the $\mathrm{3d}$ electronic transitions of the metal. For example, nickel(II) aqueous complexes $(\ce{[Ni(H2O)6]^2+})$ often absorb strongly in the red region of the visible spectrum, so the resulting color observed is green. That means the complexes with weaker ligands such as $\ce{H2O}$ absorb light with lower frequency (larger wavelengths) such as color red (see a representation of a color wheel below):


[![Color Wheel](https://i.stack.imgur.com/aKlKL.png)](https://i.stack.imgur.com/aKlKL.png)


That means, color of the solution is light blue $(\ce{[Ni(H2O)4(en)]^2+})$ means it absorbs orange color, blue $(\ce{[Ni(H2O)2(en)2]^2+})$ means it absorbs orange-yellow color, and pink (to me it looks pink, not purple) means $(\ce{[Ni(en)3]^2+})$ it absorbs yellow-green color (if it is purple, it still absorbs yellow light). If you look at these absorbed light, you'd see frequencies of yellow-green (or yellow) $\gt$ orange-yellow $\gt$ orange $\gt$ red, a good agreement with your statement.


**Note:** According to Ref.1, the $\ce{Ni^2+}$ complexes exhibit a broad band in the region $\pu{15000-23000 cm-1}$ with a second band in the range of $\pu{23000-27000 cm-1}$. These bands are assigned to the transitions $\ce{^1A\_{1g} -> ^1A\_{2g}}$ and $\ce{^1A\_{1g} -> ^1B\_{1g}}$. Also, color change due to number of ethylenediamine molecules $(\ce{H2N-CH2CH2-NH2})$ is published recently as an undergraduate experiment.




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**References:**


1. Sangamesh A. Patil and Vasant H. Kulkarni, “Complexes of nickel(II) with ethylenediamine and 1,3-diaminopropane derivatives of 2,2′-dihydroxychalkones,” *Inorganica Chimica Acta* **1983**, *73*, 125-129 (ODI: <https://doi.org/10.1016/S0020-1693(00)90836-3>).
2. Mauro Ravera, Alessandro Nucera, and Elisabetta Gabano, “Freshening up Old Methods for New Students: A Colorful Laboratory Experiment to Measure the Formation Constants of Ni(II) Complexes Containing Ethane-1,2-Diamine,” *J. Chem. Educ.* **2022**, *99(3)*, 1473–1478 (ODI: <https://doi.org/10.1021/acs.jchemed.1c01186>).


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