Question

The geometries of the ammonia complexes of ${\mathrm{Ni}}^{2+},{\mathrm{Pt}}^{2+}$ and ${\mathrm{Zn}}^{2+}$, respectively, are (A) octahedral, square planar and tetrahedral (B) square planar, octahedral and tetrahedral (C) tetrahedral, square planar and octahedral (D) octahedral, tetrahedral and square planar

Short answer

None of the provided options correctly match the geometries (tetrahedral, square planar and tetrahedral) to \(\mathrm{Ni}^{2+}\text{,} \mathrm{Pt}^{2+}\text{ and }\mathrm{Zn}^{2+}\text{ respectively. The correct geometries should be square planar for ${\mathrm{Ni}}^{2+}\text{ and }{\mathrm{Pt}}^{2+}$ and tetrahedral for \(\mathrm{Zn}^{2+}\text{), but that combination is not listed among the options.

Step-by-step solution

Determine Electron Configuration

First find the electron configuration for each metal ion. - ${\mathrm{Ni}}^{2+}$ has a d8 electron configuration. - ${\mathrm{Pt}}^{2+}$ has a d8 electron configuration. - ${\mathrm{Zn}}^{2+}$ has a d10 electron configuration.

Consider the Crystal Field Theory

Apply the crystal field theory to predict the complex geometries. - For d8 configuration in a strong field, ${\mathrm{Ni}}^{2+}$ and ${\mathrm{Pt}}^{2+}$ ions, the low spin configurations lead to a square planar geometry. - For d10 configuration like ${\mathrm{Zn}}^{2+}$, the geometry is typically tetrahedral.

Identify the Geometry for Each Complex

Based on electron configurations and the crystal field stabilization: - ${\mathrm{Ni}}^{2+}$ prefers a square planar or octahedral geometry, but in the presence of ammonia, which is a strong field ligand, square planar is preferred for a d8 ion. - ${\mathrm{Pt}}^{2+}$ also has a d8 electron configuration and, being a second and third row transition metal with strong field ligands like ammonia, typically forms a square planar complex. - ${\mathrm{Zn}}^{2+}$ has a d10 configuration and thus does not experience crystal field splitting, leading to a tetrahedral geometry, as d10 complexes do not have crystal field stabilization energy differences between different geometries.

Matching Complexes with Geometries

Match the correct geometries to the complexes based on the conclusions from Steps 1-3. - ${\mathrm{Ni}}^{2+}$ complex is square planar. - ${\mathrm{Pt}}^{2+}$ complex is also square planar. - ${\mathrm{Zn}}^{2+}$ complex is tetrahedral.

Choose the Correct Option

Since both ${\mathrm{Ni}}^{2+}$ and ${\mathrm{Pt}}^{2+}$ form square planar complexes with ammonia, and ${\mathrm{Zn}}^{2+}$ forms a tetrahedral complex, the correct answer is option (C) tetrahedral, square planar and octahedral (it appears there is a mistake in the available options, as none of them fully align with the conclusion that ${\mathrm{Pt}}^{2+}$ is square planar and ${\mathrm{Zn}}^{2+}$ is tetrahedral, we should expect ${\mathrm{Pt}}^{2+}$ to be square planar as per step 3). This indicates a correction in the options or the question might be needed.

Key concepts

Crystal Field Theory

Understanding the crystal field theory (CFT) is essential when studying the geometries of transition metal complexes. CFT explains how the arrangement of ligands around a central metal ion impacts the energy of the metal's d-orbitals. When ligands approach the metal ion, they interact with the d-orbitals that differ in energy depending on their orientation in space relative to the ligands.

For example, in an octahedral geometry, ligands approach along the axes, causing the d-orbitals to split into two energy levels: t_2g and e_g. The exact splitting pattern and energy difference, known as the crystal field splitting energy (equationDelta), are determined by the geometry of the complex and the type of the ligands. In strong-field ligands like ammonia, the splitting is significant, leading to low-spin configurations and distinct geometric preferences for certain electron configurations.

Electron Configuration

Electron configuration is fundamental to predicting the behavior of transition metal ions in complexes. Transition metals have a partially filled d-subshell, which can accommodate up to ten electrons. The way these electrons are distributed among the five d-orbitals significantly influences the properties and geometry of metal complexes.

For instance, a equationd^8 configuration can lead to square planar or octahedral geometries. However, in cases of strong-field ligands, pairing of electrons is energetically favorable, often resulting in a square planar geometry. This demonstrates how understanding electron configuration can offer insights into the structure and bonding of coordination compounds.

Transition Metal Complexes

Transition metal complexes consist of a central metal ion bonded to surrounding molecules or ions called ligands. These ligands donate electron pairs to the metal, forming coordinate covalent bonds. The combination of a metal's electron configuration and the nature of the ligands, as well as their arrangement around the metal, defines the geometry of the complex.

Transition metal complexes exhibit a wide range of geometries, such as octahedral, square planar, and tetrahedral. Each structure is characterized by the spatial distribution of ligands around the central ion, which is crucial in determining properties such as color, reactivity, and magnetic behavior. The nature of the ligands (whether they are strong or weak field) also plays a significant role in stabilizing certain geometries over others.

Ligand Field Stabilization

Ligand field stabilization is a concept that arises from the ligand field theory, an extension of crystal field theory that includes the covalent aspects of bonding in addition to the ionic interactions. It refers to the stability gained by a transition metal complex due to the specific arrangement of ligands around the central metal and the resulting electron configuration.

Stronger field ligands, which induce a larger splitting in the d-orbitals (equationDelta), can lead to lower-energy configurations by forcing electron pairing in the lower-energy orbitals before occupying the higher-energy ones, thus creating a low-spin state. This effect has significant implications on the preferred geometry of the complex, often favoring structures like square planar for low-spin d⁸ complexes due to the extra stability imparted by these arrangements.