Absorption of Light and Wavelength-Frequency Relationship in Transition Metal Complexes
What Causes Color in Transition Metal Complexes?
- Transition metal complexes are often brightly colored due to the behavior of their d-electrons.
- Transition metals have partially filled d-orbitals, and when they form complexes with ligands (molecules or ions that donate electron pairs), the normally degenerate (equal energy) d-orbitals split into two sets with different energies.
- This process, known as d-orbital splitting, happens because ligands interact more strongly with some d-orbitals than others.
- When white light shines on a transition metal complex, specific wavelengths of light are absorbed.
- This occurs because electrons in lower-energy d-orbitals absorb energy and are promoted to higher-energy d-orbitals.
- The energy difference ($\Delta E$) between these split d-orbitals matches the energy of the absorbed light.
- The wavelengths not absorbed are transmitted or reflected, determining the observed color.
- The color you observe is the complementary color of the light absorbed.
if a complex absorbs yellow light, it appears violet because violet is yellow's complementary color.
Factors Affecting Color in Transition Metal Complexes
Several factors influence the color of a transition metal complex:
- The Metal Ion: Different transition metals have varying d-electron configurations, which affect d-orbital splitting.
- The Oxidation State: A higher oxidation state increases the positive charge on the metal ion, pulling ligands closer and causing greater splitting energy.
- The Ligands: The nature of the ligands affects the splitting. Strong field ligands like cyanide ($CN^-$) cause greater splitting than weak field ligands like water ($H_2O$).
- Geometry: The spatial arrangement of ligands (e.g., octahedral, tetrahedral) alters the splitting pattern.
- The complex $[Cu(H_2O)_6]^{2+}$ appears blue because it absorbs orange light. g in
- In contrast, $[CuCl_4]^{2-}$ appears yellow-green because it absorbs violet light.
- The difference in color arises because chloride ($Cl^-$) is a weaker ligand than water, resulting in smaller d-orbital splitting.
Other examples of coloured compounds:
1. Hexaaquairon(III) Complex – Yellow-Brown
- Formula: $\text{[Fe(H}_2\text{O)}_6]^{3+}$
- Color: Yellow-brown
- Reason: Absorbs violet and blue light, transmitting yellow and red wavelengths.
2. Tetraamminecopper(II) Complex – Deep Blue (Royal Blue)
- Formula: $\text{[Cu(NH}_3\text{)}_4]^{2+}$
- Color: Deep blue
- Reason: Strong ligand field splitting causes absorption in the red-orange region, transmitting blue.
3. Hexacyanoferrate(III) Complex – Orange-Red
- Formula: $\text{[Fe(CN)}_6]^{3-}$
- Color: Orange-red
- Reason: Cyanide is a strong field ligand, causing higher energy absorption in the blue region, transmitting orange-red light.
4. Hexaamminecobalt(III) Complex – Yellow
- Formula: $\text{[Co(NH}_3\text{)}_6]^{3+}$
- Color: Yellow
- Reason: Absorbs violet and blue light, transmitting yellow.
- To predict the color of a complex, use the color wheel provided in your data booklet.
- Remember, the absorbed and observed colors are complementary.
