Why transition elements are coloured and paramagnetic

Coloured compounds

Transition metal ions have an incomplete d subshell. In the field of ligands the d orbitals split into two energy levels (t₂g and e_g). The energy gap lies in the visible region, so an electron can absorb a part of visible light and jump to the higher level (d–d transition). The remaining light is transmitted and appears as the complementary colour. Example: Cu²⁺ is blue, Fe³⁺ is yellow-brown, Mn²⁺ is pale pink. Ions with d⁰ or d¹⁰ (Sc³⁺, Zn²⁺) are colourless.

Paramagnetism

Paramagnetism is due to unpaired electrons, each acting as a small magnet. Transition metal ions have partly filled d orbitals that contain unpaired electrons (Fe³⁺ has 5, Cu²⁺ has 1), so they are attracted into a magnetic field. The magnetic moment is given by the spin-only formula μ = √(n(n+2)) BM, where n is the number of unpaired electrons. Mn²⁺ (n = 5) gives 5.92 BM.