Colour in transition metal compounds is attributed to (a) Small size metal ions. (b) Absorption of light in UV region (c) Complete (n, s) subshell (d) Incomplete (n-1) d subshell

The mesmerizing colors exhibited by transition metal compounds arise from a process called electronic transitions. Imagine these as tiny leaps taken by electrons within the d subshell of the metal ions. When light hits the compound, some photons have just the right amount of energy to boost an electron from a comfortable low energy state to an exhilarating high energy level. Because these electrons reside in the outer orbits, the energy change associated with their transitions fits perfectly into the energy of visible light.
This absorption and consequent electron jump are highly specific to the energy difference between the two levels, akin to selecting the right key for a lock. Therefore, not all light is absorbed, but rather specific wavelengths that match the energy gap. The non-absorbed light gets reflected back to our eyes, and we perceive the complementary color of the absorbed wavelength. This is why copper sulfate appears blue; it absorbs light in the red region, so we see the blue that's left over.

Transition metals have this nifty feature where they have an incomplete (n-1) d subshell. These d subshell electrons are the primary players in the color game of transition metal compounds. Unlike s or p electrons that typically fill up neatly and completely, d subshell electrons can be more flexible and accept a range of electron configurations. This is due to their ability to accommodate up to ten electrons.
Depending on how many of these seats are taken, the energy landscape within the atom changes. Some metal ions might have a nearly full d subshell, while others have just a few electrons. The various possible arrangements create different gaps between energy levels that decide which particular wavelength an electron will leap across when absorbing light. So, it's not just any metal ion that will display color; it must have that special incomplete d subshell to join the color parade.

The reason we see colors at all is thanks to our eyes' ability to detect visible light, which is a sliver of the whole electromagnetic spectrum — think of it as the rainbow range of light. When transition metal compounds absorb energy from visible light, it's this absorption that narrows down the remaining reflected or transmitted light spectrum to what we actually see.
In this context, 'visible light absorption' directly determines the perceived color. If the absorbed light doesn't fall within the visible spectrum, the compound won't contribute color in the same way. For example, ultraviolet light absorption doesn't affect visible color since we can't see UV light. This is why answer choice (b) from our exercise, which refers to UV region absorption, doesn't explain the presence of color in transition metals. It's the dance of electrons in the yet closer-to-us visible region that garners a round of applause in the form of vivid color.