Final answer:
The color of a chromium complex depends on the light it absorbs, which is influenced by its ligands. Identifying the specific ligands A and B is required to determine which complex is violet, as this color would be observed if the complex absorbed in the 560 - 570 nm range.
Step-by-step explanation:
The question is asking which of the chromium complexes will be violet in color, given that if a complex absorbs light within the 560 - 570 nm range, it appears violet. The color of a coordination complex, like [Cr(A)6]3- or [Cr(B)6]3-, is determined by the light it absorbs, which in turn depends on the ligands surrounding the central metal ion and the arrangement of its d-electrons.
Without knowing the specific ligands A and B, it is impossible to determine accurately which complex will appear violet. However, if one of these complexes has ligands that cause it to absorb light in the range of 560 - 570 nm, that complex would then reflect or transmit light at the complementary color, which is violet. For example, a complex with strong-field ligands, like [Cr(NH3)6]3+, absorbs higher energy photons, including those in the blue-violet light, leading it to appear yellow. Conversely, a complex with weak-field ligands absorbs lower energy photons, such as yellow-green light, resulting in a violet color.