Coordination Compounds – JEE Main ChemistryPractice Questions & PYQs

The oxidation state of a central metal in a coordination complex is determined by setting the sum of the charges of all ligands and the central atom equal to the total charge of the complex ion. Since ammonia is a neutral ligand and each chloride ion contributes a charge of -1, the oxidation state $x$ of the chromium atom in the given complex is found using the equation $x + 4(0) + 2(-1) = +1$. Solving this expression for the unknown oxidation state yields $x = 3$.

The magnitude of paramagnetic behavior is directly determined by the number of unpaired electrons in a substance, where a smaller quantity of unpaired electrons corresponds to lower magnetism. Among these ions, $\mathrm{Cr}^{3+}$ possesses 3 unpaired electrons, $\mathrm{Mn}^{3+}$ has 2, and $\mathrm{Fe}^{3+}$ has 1. Because the $\mathrm{Co}^{3+}$ ion in the complex $[\mathrm{Co}(\mathrm{CN})_6]^{3-}$ contains 0 unpaired electrons, it exhibits the lowest level of paramagnetic behavior.

The coordination number of a central metal ion in a complex is strictly defined as the total number of coordinate covalent bonds, or $\sigma$ bonds, that the metal has established with its surrounding ligands. Because each coordinate bond represents the donation of an electron pair from a ligand into an empty metal orbital, the coordination number is determined by summing these interactions, which involves counting each monovalent ligand as one and each bidentate ligand as two. While ligands may engage in additional interactions such as $\pi$ bonding, these do not contribute to the coordination number, which relies exclusively on the count of $\sigma$ bonds formed between the metal and the ligands.

Coordination compounds serve as essential biological agents by utilizing a central metal atom coordinated within a specific organic framework to facilitate physiological functions. Chlorophyll, the green pigment necessary for photosynthesis, contains a porphyrin ring centered around a magnesium ion rather than calcium, making the statement identifying it as a calcium-containing compound incorrect. Conversely, other biological complexes correctly feature distinct metal ions, such as zinc in carboxypeptidase-A, cobalt in cyanocobalamin ($B_{12}$), and iron in haemoglobin.

The spin-only magnetic moment of a coordination complex is determined by the number of unpaired electrons in the central metal ion, governed by the formula $\mu_{eff} = \sqrt{n(n+2)}$ B.M., where a higher value of $n$ corresponds to a larger magnetic moment. Evaluating the given complexes, $[\mathrm{MnCl}_4]^{2-}$ possesses five unpaired electrons, $[\mathrm{CoCl}_4]^{2-}$ has three unpaired electrons, and $[\mathrm{Fe}(\mathrm{CN})_6]^{4-}$ contains zero unpaired electrons due to the strong-field nature of the cyanide ligands. Consequently, the correct order based on the magnitude of magnetic moments is $[\mathrm{MnCl}_4]^{2-} \rightarrow [\mathrm{CoCl}_4]^{2-} \rightarrow [\mathrm{Fe}(\mathrm{CN})_6]^{4-}$.

To determine which coordination complex possesses the greatest variety of isomers, one must analyze the symmetry and spatial arrangement of the ligands surrounding the central metal ion. While $\mathrm{[Co(NH_3)_5Cl]^{2+}}$ exhibits no isomerism and both $\mathrm{[Ru(NH_3)_4Cl_2]^{+}}$ and $\mathrm{[Ir(PR_3)_2H(CO)]^{2+}}$ are limited to just two geometrical isomers, the complex $\mathrm{[Co(en)_2Cl_2]^{+}}$ offers greater structural diversity. Because this specific complex supports both geometrical isomerism and the existence of optical enantiomers for its cis configuration, it results in a higher total number of distinct isomers compared to the others.

The cyanide ion, $\mathrm{CN}^{-}$, possesses a lone pair of electrons on the carbon atom that allows it to function as a powerful ligand, facilitating the formation of stable coordination complexes with various metal cations. In addition to its role in complexation, this species demonstrates significant electron-donating behavior, which enables it to act effectively as a reducing agent when interacting with metal species. Because it does not typically behave as an electron acceptor in these interactions, it is characterized solely by its ability to engage in complexing and reducing processes.

Coordination complexes are classified based on whether they utilize inner $(n-1)d$ orbitals or outer $nd$ orbitals for hybridization during bond formation, where the usage of inner orbitals defines an inner orbital complex and the usage of outer orbitals defines an outer orbital complex. While the complexes $[\mathrm{Fe}(\mathrm{CN})_6]^{4-}$, $[\mathrm{Mn}(\mathrm{CN})_6]^{4-}$, and $[\mathrm{Co}(\mathrm{NH}_3)_6]^{3+}$ rely on the inner $(n-1)d$ subshell to achieve $d^2sp^3$ hybridization, $[\mathrm{Ni}(\mathrm{NH}_3)_6]^{2+}$ utilizes the outer $nd$ orbitals for bonding. This distinct utilization of the outer orbitals results in an $sp^3d^2$ configuration for $[\mathrm{Ni}(\mathrm{NH}_3)_6]^{2+}$, which classifies it as an outer orbital complex.

Coordination complexes consist of a stable coordination sphere that remains intact in solution, while ionizable groups dissociate completely to determine the total number of ions produced. Because one mole of the compound yields three moles of ions, the structure must split into one complex cation and two counter anions. The reaction with silver nitrate provides further evidence, as the precipitation of two moles of $\mathrm{AgCl}$ upon the addition of two moles of $\mathrm{AgNO}_3$ confirms that exactly two chloride ions reside outside the coordination sphere, resulting in the structure $[\mathrm{Co}(\mathrm{NH}_3)_5\mathrm{Cl}]\mathrm{Cl}_2$.

To determine the oxidation state of a central metal atom in a neutral coordination compound, we must ensure the algebraic sum of all charges equals zero. Since potassium carries a charge of +1 and each cyanide ligand carries a charge of -1, we assign $x$ to nickel and set up the following balance equation:

$4(+1) + x + (-1) \times 4 = 0$

Solving this expression yields $4 + x - 4 = 0$, which results in $x = 0$.

The formation of the $[\mathrm{Cu}(\mathrm{NH}_3)_4]^{2+}$ complex depends on the ability of ammonia to function as a Lewis base, utilizing the lone pair of electrons on its nitrogen atom to bond with copper ions. In an acidic environment, however, these lone pairs are rapidly protonated by the excess hydrogen ions present, which removes the ammonia from the system by converting it into the ammonium ion:
$\mathrm{NH}_3 + \mathrm{H}^+ \rightleftharpoons \mathrm{NH}_4^+$
Since the available ammonia molecules are sequestered as $\mathrm{NH}_4^+$, there are no longer free ligand molecules available to coordinate with the copper ion and generate the specified complex.

Organometallic compounds are defined by the presence of a direct covalent bond between a carbon atom and a metal atom. In the structure of $\mathrm{CH}_3-\mathrm{Mg}-\mathrm{Br}$, often called Grignard's reagent, the carbon atom of the methyl group is directly bonded to the magnesium metal atom. Since the molecule contains this $\mathrm{C}-\mathrm{Mg}$ linkage, it meets the essential criteria to be classified as an organometallic compound.

Linkage isomerism arises in coordination compounds that contain an ambidentate ligand capable of bonding to the central metal ion through more than one donor atom. In the case of $[\mathrm{Cr}(\mathrm{NH}_3)_5 \mathrm{NO}_2] \mathrm{Cl}_2$, the nitrite ion acts as an ambidentate ligand, allowing it to coordinate through the nitrogen atom to form the nitro isomer $[\mathrm{Cr}(\mathrm{NH}_3)_5 \mathrm{NO}_2] \mathrm{Cl}_2$ or through the oxygen atom to form the nitrito isomer $[\mathrm{Cr}(\mathrm{NH}_3)_5 \mathrm{ONO}] \mathrm{Cl}_2$. Because the ligand creates distinct structures by linking to the chromium metal center via different atoms, the compound exhibits linkage isomerism.

The stability of coordination complexes is fundamentally governed by the Lewis basicity of the participating ligands, with stronger nucleophiles forming significantly more stable bonds. Because the $\mathrm{CN}^{-}$ ion is a much more potent nucleophile compared to halides like $\mathrm{Cl}^{-}$ or neutral ligands like $\mathrm{H}_2\mathrm{O}$, it creates a more energetically favorable interaction with the central iron atom. This heightened electron-donating efficiency renders the complex $[\mathrm{Fe}(\mathrm{CN})_6]^{3-}$ the most stable among the provided options.