How do metal complexes exhibit different colors?

Metal complexes show different colors because of the way their electrons absorb and reflect light. Here’s the explanation:

Why colors appear

• In a metal complex, the metal ion is surrounded by ligands.
• The ligands affect the energy levels of the metal’s d-orbitals.
• When light shines on the complex, some wavelengths (colors) of light are absorbed to move electrons from lower-energy d-orbitals to higher-energy ones.
• The remaining light that is not absorbed is what we see as the color of the complex.

Factors that change the color

1. Type of metal ion – Different metals absorb different wavelengths.
2. Oxidation state – The same metal in different states (+2, +3, etc.) can give different colors.
   • Example: Fe²⁺ complexes (pale green) vs Fe³⁺ complexes (yellow/brown).
3. Nature of ligands – Strong ligands (like CN⁻) cause bigger energy changes than weak ligands (like H₂O).
   • This changes which light is absorbed and hence the color.
4. Geometry of the complex – Octahedral, tetrahedral, or square-planar shapes split d-orbitals differently, leading to different colors.

In short:
Metal complexes show colors because their d-electrons absorb certain parts of visible light, and the color we see is the light that remains.