Colour of complexes

Summary and Exam Tips for Colour of complexes

Colour of complexes is a subtopic of Chemistry of transition elements (A-Level Inorganic Chemistry), which falls under the subject Chemistry in the Cambridge International A Levels curriculum.

The color of transition element complexes arises from the splitting of degenerate d orbitals into non-degenerate ones when ligands form dative covalent bonds with the central metal ion. In octahedral complexes, six ligands cause the $3d_{x^2-y^2}$ and $3d_{z^2}$ orbitals to experience greater repulsion, resulting in higher energy levels compared to the $3d_{yz}$, $3d_{xz}$, and $3d_{xy}$ orbitals. Conversely, in tetrahedral complexes, the $3d_{x^2-y^2}$ and $3d_{z^2}$ orbitals are at lower energy levels due to less repulsion. The energy difference, $\Delta E$, between these orbitals is crucial for electron promotion. When light of a specific frequency is absorbed, electrons transition to higher energy levels, and the unabsorbed light frequencies combine to create the complementary color we perceive. Ligands significantly influence $\Delta E$, leading to variations in the observed color of complexes. For example, ligand exchange in copper(II) and cobalt(II) complexes results in noticeable color changes due to different $\Delta E$ values.

Exam Tips

• Understand Orbital Splitting: Be clear about how d orbitals split in octahedral and tetrahedral complexes, and how this affects the energy levels.
• Complementary Colors: Remember that the color observed is the complementary color of the absorbed light frequency.
• Ligand Influence: Know how different ligands affect the energy difference $\Delta E$ and thus the color of the complex.
• Examples: Familiarize yourself with examples like the color change in [Cu(H₂O)₆]²⁺ to [Cu(NH₃)₄(H₂O)₂]²⁺ to illustrate ligand effects.
• Equation Application: Practice using the equation $\Delta E = h \times v$ to calculate energy differences and understand electron promotion.