Question

Explain why an electron in a \(d_{z^{2}}\) or \(d_{x^{2}-y^{2}}\) orbital in an octahedral complex will experience greater repulsions because of the presence of the ligands than an electron in a \(d_{x y}, d_{x z}\), or \(d_{y z}\) orbital.

Short answer

The d_{z^{2}} and d_{x^{2}-y^{2}} orbitals in an octahedral complex are oriented directly towards the ligands, leading to greater electron-ligand repulsion compared to d_{xy}, d_{xz}, d_{yz} orbitals that point between the ligands.

Step-by-step solution

Identify the Geometry of the Complex

Recognize that the complex mentioned is octahedral, meaning it has six ligands symmetrically arranged around a central metal ion. The ligands are located along the axes in an octahedral crystal field.

Understand the Orientation of d-Orbitals

In an octahedral field, the five d-orbitals split into two sets based on energy: the higher energy set ( d_{x^{2}-y^{2}} ) and ( d_{z^{2}} ), and the lower energy set ( d_{xy}, d_{xz}, d_{yz} ). This is because the d_{x^{2}-y^{2}} and d_{z^{2}} orbitals point directly at the ligands, while the d_{xy}, d_{xz}, d_{yz} orbitals point in between the ligands.

Explain Orbital-Ligand Repulsion

Because the d_{z^{2}} and d_{x^{2}-y^{2}} orbitals are aligned towards the ligands, they experience more direct repulsion from the negative charge on the ligands. Consequently, electrons in these orbitals face more repulsion than those in the d_{xy}, d_{xz}, d_{yz} orbitals, which are oriented between the ligand axes.

Key concepts

d-orbital Repulsion

Understanding the concept of d-orbital repulsion in an octahedral complex is key to grasping the distribution of electrons around the central metal ion. In such a geometry, we imagine the metal ion at the center of an imaginary sphere with ligands positioned at equal distances along the Cartesian axes. The five d-orbitals of the central metal ion are affected differently by these ligands due to their spatial orientation.

When ligands approach the metal ion, they bring negative charge closer to the d-orbitals. However, not all d-orbitals are affected equally. The d_{z^{2}} and d_{x^{2}-y^{2}} orbitals have lobes that point directly along the axes, towards the ligands. This direct alignment causes increased electron density between the metal ion and the approaching ligands, leading to stronger electrostatic repulsions. As a result, electrons in these orbitals are pushed to a higher energy level compared to the other d-orbitals.

In contrast, the d_{xy}, d_{xz}, and d_{yz} orbitals lie at an angle to the axes, between the ligands, minimizing repulsion and allowing these orbitals to remain at a lower energy. This distinction is fundamental in predicting the electronic structure and properties of octahedral complexes.

Orbital-Ligand Interaction

The interaction between orbitals of the central metal ion and the ligands around it, known as orbital-ligand interaction, is foundational in inorganic chemistry. In an octahedral crystal field, ligands approach the central ion along the axes of a Cartesian coordinate system and interact with the d-orbitals.

This interaction is not uniform as it varies based on the spatial orientation of each d-orbital. The extent of interaction is greater for orbitals with lobes or electron density pointing directly at the ligands. This is precisely what happens to the d_{z^{2}} and d_{x^{2}-y^{2}} orbitals which experience direct overlap with the approaching ligands' orbitals. This enhanced interaction results in a higher repulsion and therefore an increased energy level for electrons found within these orbitals.

Conversely, the remaining d-orbitals, d_{xy}, d_{xz}, and d_{yz}, point between the incoming ligand axes. The reduced overlap with the ligands means less repulsion and energy stabilization for electrons in these orbitals. These interactions explain the concept of crystal field splitting, a pivotal aspect of the crystal field theory used to describe the electronic structures of transition metal complexes.

Octahedral Complex Geometry

When considering the octahedral complex geometry, envision a central metal ion surrounded at equal distances by six ligands positioned at the corners of an octahedron. This three-dimensional shape is one of the most common geometries for coordination compounds. Its implications for the behavior of d-orbitals are greatly consequential for the electronic structure and chemical properties of the complex.

The symmetrical shape of the octahedron directs the ligands along the x, y, and z axes, providing an ideal framework to analyze the effects on the five d-orbitals of the central ion, leading to the concept of crystal field splitting. Electrons within the metal ion's d-orbitals will fill the lower energy levels first, as per the Aufbau principle, and these relative energies help determine properties such as magnetism and color.

Understanding the subtleties of octahedral geometry forms the basis for comprehending more complex molecular geometries and is essential in predicting the reactivities and interactions of coordination compounds in chemical reactions.