91Ó°ÊÓ

In coordination chemistry, oxidation state is crucial for determining the electron count of a metal complex. It helps in predicting the chemical reactivity and stability of the compound. Oxidation state of a metal is the charge left on the metal atom after all the ligands are removed along with their associated electrons.
For complex (A) \([\mathrm{Co}\left(\mathrm{NH}_{3}\right)_{4}\left(\mathrm{H}_{2}\mathrm{O}\right)_{2}\right] \mathrm{Cl}_{2}\), cobalt has an oxidation state of \(+2\). We know this because the complex has a neutral charge. Each chloride ion contributes a \(-1\) charge, totaling \(-2\) for two chlorides. The ammonia and water ligands are neutral. Thus, cobalt must be \(+2\) to balance the charges.

• (A) \[\text{Oxidation State: Co}^{2+}\]

Similarly, in (B) \([\mathrm{Pt}\left(\mathrm{NH}_{3}\right)_{2} \mathrm{Cl}_{2}]\), platinum is in the \(+2\) oxidation state. Ammonia is neutral and each chloride contributes \(-1\), so the net charge on Pt is \(+2\) to balance with the two chlorides.

• (B) \[\text{Oxidation State: Pt}^{2+}\]

For complex (C) \([\mathrm{Co}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5}\mathrm{Cl}]\mathrm{Cl}\), cobalt is in the \(+3\) oxidation state. The external chloride is \(-1\), leaving cobalt as \(+3\) to maintain the compound as neutral.

• (C) \[\text{Oxidation State: Co}^{3+}\]

Finally, for (D) \([\mathrm{Ni}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}]\mathrm{Cl}_{2}\), nickel is in the \(+2\) oxidation state. The two chlorides provide \(-2\) total charge, so nickel must be \(+2\) to neutral the overall charge.

• (D) \[\text{Oxidation State: Ni}^{2+}\]

Knowing the magnetic properties of a complex can tell us about the arrangement of electrons and the nature of the bonding. Complexes can be paramagnetic, indicating unpaired electrons, or diamagnetic, indicating all electrons are paired.
Complexes with unpaired electrons exhibit magnetism when placed in a magnetic field. This is because each unpaired electron possesses a magnetic moment. In comparison, diamagnetic complexes are those with paired electrons and do not exhibit magnetism.

• Complex (A) \([\text{Co(NH}_3\text{)}_4(\text{H}_2\text{O)}_2]\text{Cl}_2\) is paramagnetic. Cobalt in the \(+2\) oxidation state has unpaired electrons.


• Complex (B) \([\text{Pt(NH}_3\text{)}_2\text{Cl}_2]\) is diamagnetic. Pt in the \(+2\) oxidation state often forms a square planar geometry, usually leading to paired electrons.


• Complex (C) \([\text{Co(H}_2\text{O)}_{5}\text{Cl}]\text{Cl}\) is paramagnetic due to the \(+3\) oxidation state of Co, which introduces unpaired electrons in the d-orbitals.


• Complex (D) \([\text{Ni(H}_2\text{O)}_6]\text{Cl}_2\) is paramagnetic. Ni \(+2\) also encounters unpaired electrons in its d-orbitals.

Geometrical isomerism is an important concept in coordination chemistry and occurs in complexes where the arrangement of ligands around the metal center allows different spatial arrangements. It happens commonly in square planar and some octahedral complexes.
In complex (B) \([\mathrm{Pt}\left(\mathrm{NH}_{3}\right)_{2} \mathrm{Cl}_{2}]\), geometrical isomerism is possible due to its square planar shape. Here, the ammonia and chloride ligands can be arranged differently around the platinum center, leading to different isomers:

• Cis isomer: Where similar ligands, say \(\mathrm{NH}_3\), are adjacent to each other.


• Trans isomer: Where similar ligands are opposite each other across the metal center.

The remaining complexes [(A), (C), and (D)] do not exhibit geometrical isomerism due to their respective geometries (octahedral for A, C, and D) where such isomerism isn’t possible given the type and arrangement of ligands.