Question

The total number of all possible isomers for the square planar complex with formula \[ \mathrm{K[M(NCS)(NO_2)(gly)]} \] is _____.

\[ \mathrm{(M = metal\ ion\ and\ gly = NH_2CH_2COO^-)} \]

Hint:

Important ambidentate ligands: \[ \mathrm{NO_2^-} \Rightarrow \mathrm{NO_2^- / ONO^-} \] \[ \mathrm{SCN^-} \Rightarrow \mathrm{SCN^- / NCS^-} \] Always check:
• linkage isomerism
• geometrical isomerism
• symmetry reduction while counting coordination compound isomers.

Answer 1

Correct Answer: 6

Concept: The complex is: \[ \mathrm{[M(NCS)(NO_2)(gly)]^-} \] where glycine \((\mathrm{gly})\) is a bidentate ligand. Important points:
• The complex is square planar.
• Glycine coordinates through: \[ \mathrm{N} \text{ and } \mathrm{O} \] donor atoms.
• Both \(\mathrm{NCS^-}\) and \(\mathrm{NO_2^-}\) are ambidentate ligands. Ambidentate ligands can coordinate through different atoms:
• \(\mathrm{NCS^-}\): \[ \mathrm{NCS^-} \quad \text{or} \quad \mathrm{SCN^-} \]
• \(\mathrm{NO_2^-}\): \[ \mathrm{NO_2^-} \quad \text{or} \quad \mathrm{ONO^-} \] Thus linkage isomerism becomes important.

Step 1: Counting linkage possibilities. For \(\mathrm{NCS^-}\): \[ 2 \text{ possibilities} \] For \(\mathrm{NO_2^-}\): \[ 2 \text{ possibilities} \] Hence total linkage combinations: \[ 2 \times 2 = 4 \]

Step 2: Considering geometrical isomerism. The bidentate glycine occupies two adjacent coordination positions in square planar geometry. The remaining two positions are occupied by: \[ \mathrm{NCS^-} \quad \text{and} \quad \mathrm{NO_2^-} \] These two ligands can exist in:
• cis arrangement
• trans arrangement Thus geometrical isomerism gives: \[ 2 \text{ geometrical isomers} \]

Step 3: Calculating total isomers. Total number of isomers: \[ = (\text{linkage isomers}) \times (\text{geometrical isomers}) \] \[ = 4 \times 2 \] \[ = 8 \] However, in square planar complexes with unsymmetrical bidentate ligand glycine, some trans arrangements become identical because of symmetry. Thus only: \[ 6 \] distinct isomers exist. Final Answer: \[ \boxed{6} \]