Why transition metal ions are coloured

Most transition metal ions are coloured because of d–d transitions.

1. In a compound or solution the ion is surrounded by ligands (water molecules), which split the five degenerate d orbitals into two sets of slightly different energy (t₂g and e_g in an octahedral field).
2. If the d subshell is partly filled, an electron can absorb a photon of visible light and jump from the lower (t₂g) to the higher (e_g) level. The energy gap Δ is of the order of visible light.
3. The ion therefore absorbs some wavelengths from white light; the remaining complementary colour is transmitted and seen.

 e_g  ─── ───            ↑  electron excited
        ↑ hν (visible)   │
 t₂g  ─── ─── ───        │ Δ
 (d–d transition: absorbed colour → complementary colour seen)

Examples: [Cu(H₂O)₄]²⁺ blue, [Ni(H₂O)₆]²⁺ green, [Ti(H₂O)₆]³⁺ purple, [Cr(H₂O)₆]³⁺ violet, [Co(H₂O)₆]²⁺ pink, [Mn(H₂O)₆]²⁺ pale pink.

Ions with d⁰ (Sc³⁺, Ti⁴⁺) or d¹⁰ (Zn²⁺, Cu⁺) have no d–d transition and are colourless. Ions such as MnO₄⁻ and CrO₄²⁻ are coloured by charge transfer.