Question

The d-orbitals participating in hybridization of central metal atom may be from the outermost shell or the penultimate shell. This depends on the nature of metal and the nature of ligand. The complexes involving the inner \(\mathrm{d}\) level (inner orbital complexes) result when the ligand is a powerful or strong ligand resulting in diamagnetic or low spin complexes. A weak ligand usually results in the formation of outer orbital complex or high spin complex. The hybridization of \(\mathrm{Cu}\) in \(\left[\mathrm{CuCl}_{4}\right]^{2-}\) and \(\left[\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+}\) are respectively. (a) \(\mathrm{sp}^{3}, \mathrm{dsp}^{2}\) (b) \(\mathrm{dsp}^{2}, \mathrm{dsp}^{2}\) (c) \(\mathrm{dsp}^{2}, \mathrm{sp}^{3}\) (d) \(\mathrm{sp}^{3}, \mathrm{sp}^{3}\)

Short answer

(a) \(\text{sp}^3\), \(\text{dsp}^2\)

Step-by-step solution

Determine the Electronic Configuration of Copper

Copper has an atomic number of 29, so its electron configuration is \([\text{Ar}] \: 3d^{10} 4s^{1}\). When forming complexes, Copper often loses its single \(4s\) electron first, making it \([\text{Ar}] \: 3d^{9}\).

Analyze the Complex [CuCl4]^{2-}

In \([\text{CuCl}_4]^{2-}\), chloride (\(\text{Cl}^-\)) is a weak field ligand. Weak ligands do not cause pairing of electrons in the \(3d\) orbitals. As a result, Copper in this complex will use its \(4s\), \(4p\), and the remaining \(4d\) orbitals to form an \(\text{sp}^3\) hybridized state for tetrahedral geometry.

Analyze the Complex [Cu(NH3)4]^{2+}

In \([\text{Cu}(\text{NH}_3)_4]^{2+}\), ammonia (\(\text{NH}_3\)) is a strong field ligand. Strong ligands cause electrons in the \(3d\) orbitals to pair up. Copper uses its \(3d\) and \(4s\) orbitals, in addition to some \(4p\) orbitals, to form a \(\text{dsp}^2\) hybridized state, resulting in a square planar geometry due to electron pairing.

Key concepts

Inner and Outer d-Orbitals

In chemistry, the concept of inner and outer d-orbitals is crucial for understanding hybridization in coordination complexes. These orbitals are derived from the electron configuration of transition metals, which have incomplete d-subshells.

• Inner d-orbitals: These are orbitals from the penultimate (second to last) shell. They are sometimes used in coordination with strong field ligands. This results in complexes where the d-electrons pair up, making room for ligand electrons. These complexes tend to be low spin or diamagnetic.
• Outer d-orbitals: These orbitals are from the outermost shell of the metal ion. Weak field ligands interact with these d-electrons causing minimal electron pairing, leading to high spin or paramagnetic complexes.

The key takeaway is that inner orbital complexes form low spin states, and outer orbital complexes form high spin states.

Strong and Weak Field Ligands

Strong and weak field ligands play a major role in defining the properties of coordination complexes. These terms refer to the ability of a ligand to induce electron pairing in a metal's d-orbitals.

• Strong field ligands: These ligands cause a large splitting of the d-orbitals. They promote electron pairing within the d-orbitals and lead to low spin configurations. The common strong field ligands include CN⁻, NH₃, and en (ethylene diamine).
• Weak field ligands: These lead to smaller d-orbital splitting and minimal electron pairing, resulting in high spin complexes. Typical weak field ligands include halides like Cl⁻, F⁻, and others such as H₂O.

Understanding ligand field strength helps predict the geometry and spin state of the metal complex.

Tetrahedral and Square Planar Geometries

Coordination complexes can adopt various geometries, with tetrahedral and square planar being two prevalent forms influenced by the type of ligand and metal ion involved.

• Tetrahedral geometry: Commonly associated with weak field ligands leading to high spin states, the tetrahedral arrangement involves four ligand atoms forming a shape like a triangular pyramid. It's less effective in causing electron pairing, seen in complexes like \([\text{CuCl}_4]^{2-}\).
• Square planar geometry: Typically found with strong field ligands, this geometry involves four ligand atoms arranged at the corners of a square plane around the central metal, promoting electron pairing and resulting in low spin states. This is the case for compounds such as \([\text{Cu}(\text{NH}_3)_4]^{2+}\).

The geometry of a complex affects its electronic properties and reactivity, underscoring the importance of understanding these spatial arrangements.