The color of the transition metal complex solution is dependent on: the metal, the metal oxidation state, and the number of metal d-electrons. For example iron(II) complexes are green and iron(III) complexes are orange/brown.2 Charge Transfer Bands If color is dependent on d-d transitions, why is it that some transition metal complexes are intensely colored in solution but possess no d electrons? In transition metal complexes a change in electron distribution between the metal and a ligand give rise to charge transfer (CT) bands.1 CT absorptions in the UV/Vis region are intense (ε values of 50,000 L mole-1 cm-1 or greater) and selection rule allowed. The intensity of the color is due to the fact that there is a high probability of these transitions taking place. Selection rule forbidden d-d transitions result in weak absorptions. For example octahedral complexes give ε values of 20 L mol-1 cm-1 or less.2 A charge transfer transition can be regarded as an internal oxidation-reduction process. 2 Ligand to Metal and Metal to Ligand Charge Transfer Bands Ligands possess \(σ\), \(σ^*\), \(π\), \(π^*\), and nonbonding (\(n\)) molecular orbitals.