Question

Which of the following complex ions is (are) square-planar? (a) \(\left[\mathrm{Ti}(\mathrm{CN})_{4}\right]^{2}\) (c) \(\left[\mathrm{Zn}(\mathrm{CN})_{4}\right]^{2}\) (b) \(\left.| \mathrm{Ni}(\mathrm{CN})_{4}\right]^{2-}\) \((\mathrm{d})\left[\mathrm{Pt}(\mathrm{CN})_{4}\right]^{2-}\)

Short answer

[Ni(CN)4]2- and [Pt(CN)4]2- are square-planar.

Step-by-step solution

Understand the criteria for square-planar geometry

Square-planar complexes usually occur with central atoms that have a 16-electron configuration and are typically d^8 metal complexes. Common examples include Ni(II), Pd(II), and Pt(II) with strong field ligands like CN^-.

Analyze each option

Evaluate each of the given complex ions to determine if they follow the 16-electron rule and possess d^8 configuration, suggesting they may adopt a square-planar geometry.

Evaluate [Ti(CN)4]2-

Ti in [Ti(CN)4]2- is Ti^2+. Titanium usually exhibits d^0 configuration in this oxidation state and, therefore, unlikely to form a square-planar complex.

Evaluate [Zn(CN)4]2-

Zn in [Zn(CN)4]2- is Zn^2+. This leads to a d^10 configuration as zinc has a completely filled d-subshell, typically forming tetrahedral or trigonal bipyramidal complexes rather than square-planar.

Evaluate [Ni(CN)4]2-

Ni in [Ni(CN)4]2- is known as Ni^2+. In this oxidation state, nickel has a d^8 configuration, which is characteristic of square-planar complexes, especially with strong field ligands like CN^-.

Evaluate [Pt(CN)4]2-

Pt in [Pt(CN)4]2- is Pt^2+. Platinum also has a d^8 electron configuration, which often results in a square-planar geometry with strong ligands like CN^-.

Key concepts

Complex Ions

Complex ions are fascinating structures that form when central metal atoms bind to surrounding molecules or ions called ligands. These entities are crucial in many chemical processes and have specific geometric shapes. The nature and arrangement of these ligands around the central metal atom can greatly influence the properties and reactivity of the complex.
In chemistry, complex ions can vary widely in structure, including octahedral, tetrahedral, and square-planar geometries. The geometry depends largely on factors such as the type of metal, its oxidation state, and the strength and type of ligands involved.
When discussing complexes, we often refer to them as either strong field or weak field based on the types of ligands. For instance, ligands like cyanide ( ext{CN}^-), which are considered strong field ligands, often lead to specific geometric arrangements, like square-planar structures in certain complexes.

Electron Configuration

The electron configuration of an atom, especially within complex ions, is crucial for understanding its chemical behavior and geometry. It dictates how electrons are distributed across the orbitals of an atom, influencing properties such as magnetism and color.
Transition metals, often at the center of complex ions, have unique electron configurations. This uniqueness arises from their electron distribution across the ext{d}-orbitals, which can further be altered by the formation of complexes with specific ligands.
A ext{d}^8 electron configuration is particularly noteworthy. This occurs when a transition metal has 8 electrons filling its d-orbitals. Metals with this configuration, especially when paired with strong field ligands like ext{CN}^-, tend to form square-planar complexes. Renowned examples include nickel, palladium, and platinum when they exhibit a +2 oxidation state.

Transition Metals

Transition metals form the backbone of many complex ions due to their ability to accommodate various electron configurations, bond with a wide range of ligands, and exist in multiple oxidation states. These metals are typically found in the d-block of the periodic table.
The characteristics of transition metals allow them to form different coordination geometries, including tetrahedral, octahedral, and square-planar. These geometries are determined by both the metal's electron configuration and the nature of the ligands it bonds with.
For instance, transition metals like nickel ( ext{Ni}^{2+}), palladium ( ext{Pd}^{2+}), and platinum ( ext{Pt}^{2+}) often form square-planar complexes when bonded with strong field ligands such as ext{CN}^-. This geometry minimizes repulsion and stabilizes the complex with a unique electronic distribution.

Ligand Field Theory

Ligand Field Theory (LFT) provides insight into why complex ions adopt specific shapes. It focuses on the interaction between metal atoms and the ligands around them - mainly explaining the effect these ligands have on the metallic d-orbitals.
According to LFT, when ligands approach a central metal ion, they create an electrostatic field that alters the energies of its d-orbitals. Strong field ligands such as ext{CN}^- can cause a significant split in the d-orbital energies, leading to unique configurations such as square-planar arrangements, particularly in metals with a ext{d}^8 electron count.
The energy level splitting determines whether the complex will be high spin or low spin, affecting properties like magnetism. Square-planar complexes often represent the low-spin state with maximum electron pairing in strong field environments, creating a stable and distinct electronic arrangement.