The solutions of (i) [Cr(OH2)6]3+ ions are pale blue-green, but the chromite ion (ii) [CrO4]2- is an intense yellow due to :
d-d transition in (i) and charge transfer transition in (ii)
The clue is the intensity contrast: one solution is pale, the other intense. That difference is diagnostic of the type of electronic transition responsible.
(i) [Cr(OH2)6]3+ — d-d transition. Chromium is Cr(III), a d3 ion, so partly filled d orbitals are available and the colour arises from promoting an electron within the d manifold. In a centrosymmetric octahedral complex such transitions are Laporte forbidden (g → g) and also spin constrained, so molar absorptivities are low, typically only a few to a few tens of dm3 mol-1 cm-1. Weak absorption gives a pale colour, exactly as observed.
(ii) [CrO4]2- — charge transfer transition. Here chromium is Cr(VI), which is d0. With no d electrons at all, a d-d transition is impossible, so the colour cannot come from that origin. Instead an electron is promoted from an oxygen lone-pair orbital to an empty metal d orbital — a ligand-to-metal charge transfer (LMCT). Such transitions are both Laporte and spin allowed, giving molar absorptivities of the order of 103-104, hence the intense yellow.
The d0 configuration is the decisive test: it rules out options that assign a d-d transition to the chromate ion, and the pale colour of the aqua ion rules out options that assign charge transfer to it.
Hence the colours are due to a d-d transition in (i) and a charge transfer transition in (ii).
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