Transition metals form coloured ions or compounds due to the partially filled d-orbitals. In the presence of solvent molecules (in solutions) or ligands (in complexes) or counter ions (in crystals), the d-orbitals split into two sets. The electrons in transition metal ions which occupy one set of d-orbitals having lower energy can be excited to another set of d-orbitals having high energy by absorbing energy from visible light. Since the energy difference \((\Delta E)\) is small between the two sets of d-orbitals, the light in the visible region only is absorbed by the electron during its excitation. The colour of the transition metal ion is due to d–d excitation or d–d transition of the electron. During d–d excitation the electron absorb one colour in the visible light and thus it appears in the complimentary colour of the absorbed light. The complimentary colours can be identified using Munsell colour wheel (as depicted below).

The number of electrons undergoing d–d transition and the energy difference between the two sets of orbitals decide the colour. The colour of particular transition metal ion, e.g., copper ion which is blue in aqueous solution changes to dark blue in the presence of sufficient ammonia and to green if sufficient chloride ions are added. \(\Delta\)E value depends on the nature of metal ion, the nature of ligands and several other factors. Some metal ions exhibit different colours in different oxidation states.

Which of the following are coloured though they do not contain unpaired electrons?

All of them

The correct answer is option 4. all of them.

The color of a compound arises from the interaction of light with the electrons within its molecular or ionic structure. Specifically, the color is determined by the absorption of certain wavelengths of light by the compound. Even though a compound may not contain unpaired electrons, it can still exhibit color due to various electronic transitions within its structure. Let us explore the explanations for each of the compounds mentioned:

1. \(K_2CrO_4\) (Potassium chromate):

\(K_2CrO_4\) appears yellow in color. This color arises due to electronic transitions involving the chromium ions (Cr⁶⁺). In this compound, the yellow color is attributed to the absorption of certain wavelengths of light by the Cr⁶⁺ ions.

While chromium in its +6 oxidation state typically suggests the presence of unpaired electrons in its d orbitals, in the case of \(K_2CrO_4\), the color is primarily due to d-d transitions rather than the presence of unpaired electrons.

2. \(KMnO_4\) (Potassium permanganate):

\(KMnO_4\) appears purple in color. Similar to \(K_2CrO_4\), the color arises due to electronic transitions involving the manganese ions (Mn⁷⁺). The purple color is attributed to the absorption of specific wavelengths of light by the Mn⁷⁺ ions. Again, despite the high oxidation state of manganese, the color of \(KMnO_4\) is primarily due to d-d transitions rather than the presence of unpaired electrons.

3. \(CuO_2\) (Copper(II) oxide):

\(CuO_2\) is a hypothetical compound, and it seems there might be confusion regarding its identity. Copper(II) oxide (CuO) is black in color. The color of \(CuO\) arises due to electronic transitions involving the copper ions (Cu²⁺). These transitions may include charge transfer transitions or other mechanisms.

Despite not containing unpaired electrons, \(CuO\) exhibits color due to the absorption of specific wavelengths of light by the Cu²⁺ ions, which leads to the perception of a black color.

In summary, even though these compounds do not contain unpaired electrons, they still exhibit color due to electronic transitions involving the metal ions within their structures. These electronic transitions may include d-d transitions, charge transfer transitions, or other mechanisms, depending on the specific compound and its molecular or ionic structure.