Colour of complexes — common mistakes

Remember the key relationship: Stronger field ligand → Larger ΔE → Higher frequency (shorter wavelength) light absorbed. This is a common source of exam questions, especially when explaining colour changes during ligand substitution reactions.

A common trap involves copper. Copper(II) compounds (Cu²⁺, d⁹) are almost always coloured (usually blue or green). However, copper(I) compounds (Cu⁺, d¹⁰) are colourless. Pay close attention to the oxidation state.

No. The pattern described (t₂g lower, e_g higher) is for octahedral complexes. For tetrahedral complexes, the splitting pattern is inverted: the e_g set is lower in energy and the t₂g set is higher. This is because the ligands approach 'between' the axes, repelling the d_xy, d_xz, and d_yz orbitals more. The splitting energy (ΔE_tet) is also significantly smaller than for an equivalent octahedral complex.

White light is a mixture of all colours (like a rainbow). When it passes through the complex solution, certain colours (wavelengths) are absorbed by the electrons for d-d transitions. The colours that are not absorbed pass through the solution and reach our eyes. Our brain processes this combination of remaining colours and perceives it as a single colour, which is the complement of what was removed. For example, removing orange from white light leaves behind all the other colours, which combine to look blue.

Yes. A well-known example is the permanganate ion, MnO₄⁻, which is intensely purple. Here, the manganese is in a +7 oxidation state (d⁰), so d-d transitions are impossible. The colour arises from a different mechanism called a 'charge-transfer transition', where an electron moves from an orbital that is primarily on the ligand (oxygen) to an empty orbital on the metal ion. This is beyond the scope of the A-level syllabus but explains why some d⁰ complexes are still coloured.