MCQ Detailed View

Transition metal compounds are well known for their bright colours. The origin of these colours is linked to the electronic structure of transition elements. Transition metals have partially filled (n–1) d-orbitals, also called d-subshells. When these ions form complexes with ligands, the d-orbitals split into two sets of different energy levels due to the ligand field. This phenomenon is described by Crystal Field Theory (CFT).

When visible light falls on a transition metal compound, electrons in the lower energy d-orbitals absorb a specific wavelength of light and get excited to higher energy d-orbitals. This process is called a d–d electronic transition. The colour observed is the complementary colour of the light absorbed. This is why many complexes of transition metals, such as [Cu(H₂O)₆]²⁺ (blue), [Ni(H₂O)₆]²⁺ (green), and [Cr(H₂O)₆]³⁺ (violet), appear coloured.

If the (n–1) d-orbitals are completely filled, as in Zn²⁺ (3d¹⁰), Cd²⁺ (4d¹⁰), and Hg²⁺ (5d¹⁰), no d–d transitions are possible. These ions form colourless compounds. Similarly, if the d-subshell is empty, as in Sc³⁺ (3d⁰) or Ti⁴⁺ (3d⁰), the compounds are also colourless because no electrons are available for excitation.

The incorrect options can be explained:

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  Small-sized metal ions are not responsible for colour.

  
  


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  Absorption of UV light usually leads to colourless compounds since UV is not in the visible range.

  
  


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  A complete ns sub-shell has no direct role in colour.

  
  

Thus, the main factor responsible for the colour of transition metal compounds is the presence of an incomplete (n–1) d sub-shell, which allows visible light absorption and d–d transitions. This principle is one of the most important characteristics of transition elements studied in inorganic chemistry.