Coordination Compounds – JEE Main ChemistryPractice Questions & PYQs

Ambidentate ligands are defined as species that possess at least two distinct potential donor atoms, although they coordinate to the central metal atom through only one of these sites at any specific time. Examining the provided list, $\mathrm{NO}_2^{-}$ can bond via nitrogen or oxygen, $\mathrm{SCN}^{-}$ can bond through sulfur or nitrogen, and $\mathrm{CN}^{-}$ can utilize either the carbon or nitrogen atom for bonding. Since the remaining species in the set do not exhibit this specific dual-binding capability, there are exactly 3 ambidentate ligands present.

The coordination geometry of a metal complex is primarily dictated by the total number of atoms or ligands directly bonded to the central metal center. In $\mathrm{Mn}_2(\mathrm{CO})_{10}$, each manganese atom is coordinated to five carbonyl ligands and is connected to the second manganese atom through a single metal-metal bond, establishing a total coordination number of six. This coordination environment forces the ligands into an octahedral arrangement around each manganese center.

Catalytic industrial processes rely on highly specific reagents to drive chemical transformations effectively, making the accurate identification of catalyst-process pairs essential. While silver bromide is standard for photography, the Haber process utilizes iron, and the polymerization of ethylene into polythene is achieved using a Ziegler-Natta catalyst consisting of $\mathrm{TiCl}_4$ and $\mathrm{Al}(\mathrm{CH}_3)_3$. The Wacker process, however, relies specifically on $\mathrm{PdCl}_2$ to facilitate the oxidation of ethylene, rendering the association with $\mathrm{PtCl}_2$ chemically incorrect.

The hybridization state of a coordination complex determines its spatial geometry, and in this case, $\text{sp}^3$ hybridization indicates a tetrahedral arrangement. Within a tetrahedral molecule, every ligand occupies a position that is equivalent to all others, maintaining a consistent bond angle relative to every other ligand regardless of orientation. Because no two ligands can occupy positions that are distinct from one another, tetrahedral complexes such as $MABXL$ cannot support the existence of geometrical isomers.

Iron(III) possesses a $3d^5$ configuration, and the observation that the $e_g$ orbital is empty characterizes the species as a low-spin octahedral complex where strong ligands force all five electrons to pair within the $t_{2g}$ orbitals. Since octahedral complexes involve six coordination sites, the combined number of ammonia and cyanide ligands must total 6. By assessing the overall complex charge of $-1$, the relationship $3 + x(0) + y(-1) = -1$ dictates that $y$ equals 4. With a total coordination number of 6, it follows that $x$ must be 2, resulting in a sum of $x+y=6$.

Among the given elements, Chromium corresponds to the metal with the smallest metallic radius. Within the $\mathrm{CrO}_4^{2-}$ ion, the oxidation state of Chromium is determined to be $+6$, which leaves the metal ion with an electron configuration of $d^0$. Given that the spin-only magnetic moment is calculated as $\mu = \sqrt{n(n+2)}$ BM, where $n$ is the number of unpaired electrons, an absence of unpaired electrons results in a magnetic moment of $0$ BM.

Both $\mathrm{N}(\mathrm{CH}_3)_3$ and $\mathrm{P}(\mathrm{CH}_3)_3$ possess a lone pair of electrons, which qualifies them as Lewis bases capable of donating electron density to coordinate with transition metals and form complexes. While they share this primary donor ability, their bonding behavior diverges significantly because phosphorus possesses empty d-orbitals that allow it to exhibit $\pi$-acceptor character, a property that is completely absent in nitrogen. Consequently, despite both elements belonging to the same group of the periodic table, the nature of their chemical bonding with transition metals is fundamentally different.

Hybridization is determined by the central atom's coordination geometry and whether it utilizes inner or outer $d$-orbitals for bonding, often influenced by the field strength of the ligands. In $[\mathrm{Co}(\mathrm{NH}_3)_6]^{3+}$, the presence of strong field ammonia ligands forces the pairing of electrons within the central Cobalt ion, allowing it to utilize two inner $d$ orbitals to form an octahedral $d^2sp^3$ complex. In contrast, molecules such as $\mathrm{BrF}_5$ and $\mathrm{SF}_6$ involve outer $sp^3d^2$ hybridization, whereas $[\mathrm{PtCl}_4]^{2-}$ exhibits $dsp^2$ hybridization consistent with its square planar geometry.

In a tetrahedral crystal field, the $d$-orbitals split into a lower energy $e$ set and a higher energy $t_2$ set, where electronic filling depends on the oxidation state and $d$-electron count of the central metal ion, with weak-field ligands like chloride promoting high-spin configurations. For $\mathrm{TiCl}_4$, the titanium ion is in the $+4$ oxidation state with a $3d^0$ configuration, matching the $e^0, t_2^0$ configuration. In $[\mathrm{FeO}_4]^{2-}$, iron exists in the $+6$ oxidation state, which possesses $2$ $d$-electrons that fill the $e$ orbitals to form $e^2, t_2^0$. The complex $[\mathrm{FeCl}_4]^{-}$ contains $Fe^{3+}$ with a $3d^5$ configuration, resulting in high-spin filling of $e^2, t_2^3$. Finally, $[\mathrm{CoCl}_4]^{2-}$ contains $Co^{2+}$ with a $3d^7$ configuration, which distributes electrons as $e^4, t_2^3$. Matching these results provides the sequence (A)-(III), (B)-(I), (C)-(IV), and (D)-(II).

The complex $\left[\mathrm{Pt}(\mathrm{NH}_3)_{2}\mathrm{Cl}(\mathrm{NH}_2\mathrm{CH}_3)\right]\mathrm{Cl}$ contains platinum in the $+2$ oxidation state, balancing the presence of the chloride counter-ion. As a $5d$ transition metal with an atomic number of $78$, platinum in this $+2$ state adopts a $5d^8$ electronic configuration. The strong crystal field splitting typical of such heavy transition metals forces electrons to pair completely within the $d$-orbitals, facilitating $dsp^2$ hybridization for square planar geometry. With no unpaired electrons remaining, the spin-only magnetic moment calculated using $\mu = \sqrt{n(n+2)}$ results in 0 B.M.

The magnetic property of a transition metal ion is primarily dictated by the number of unpaired electrons in its d-orbitals, which relates to the spin-only magnetic moment through the formula $\mu = \sqrt{n(n+2)} \,\mathrm{BM}$. Given a magnetic moment of $3.86\,\mathrm{BM}$, we equate this to the formula, resulting in $\sqrt{n(n+2)} \approx 3.86$, which simplifies to $n^2 + 2n \approx 15$. Solving this quadratic equation indicates that the ion must contain exactly $n = 3$ unpaired electrons.

Reviewing the configurations of the given ions, ${}_{22}\mathrm{Ti}^{+2}$ adopts a $3d^2$ arrangement with two unpaired electrons, while ${}_{23}\mathrm{V}^{+2}$ exhibits a $3d^3$ configuration, which perfectly matches the requirement of having three unpaired electrons. Comparing this to the other options, ${}_{25}\mathrm{Mn}^{+2}$ has a $3d^5$ configuration with five unpaired electrons and ${}_{26}\mathrm{Fe}^{+2}$ has a $3d^6$ configuration with four, confirming that the metal with atomic number 23 is the only one consistent with the observed magnetic moment.

The spin state of a coordination complex depends on the strength of the ligand field, which dictates whether electrons within the metal's $d$-orbitals will pair up or remain unpaired. In ${[\mathrm{Co}(\mathrm{NH}_3)_6]^{3+}}$, the strong field ammonia ligand forces the $3d^6$ electrons of the ${Co^{3+}}$ ion to pair into the $t_{2g}^6 e_g^0$ configuration with $d^2sp^3$ hybridization, characterizing it as a spin-paired complex. In contrast, the weak field fluoride ligand in ${[\mathrm{CoF}_6]^{3-}}$ results in an $sp^3d^2$ hybridized structure with a $t_{2g}^4 e_g^2$ configuration where electrons remain unpaired, classifying it as a spin-free complex.

The splitting of d-orbitals in a coordination complex is directly determined by the strength of the coordinating ligands, where stronger ligands create a larger energy gap between the split orbitals. Because the energy of the photon required for electronic transition, expressed by the relationship $E = hc\bar{\nu}$, is directly proportional to the wavenumber, complexes with larger field splitting absorb light at higher wavenumbers. Consequently, the correct sequence is derived by simply ranking the complexes based on their ligand field strength, as the increasing magnitude of this splitting directly dictates the rise in the wavenumber of the absorbed light.

Homoleptic complexes are defined by a coordination sphere where the central metal atom is bonded exclusively to a single type of ligand, indicating complete uniformity in the surrounding chemical environment. Within the provided list, $[\mathrm{Ni}(\mathrm{CN})_4]^{2-}$ qualifies as homoleptic because all its attached ligands are cyanide, whereas the other candidates feature a mixture of different ligands like ammonia and chloride. This consistent attachment of identical groups to the metal center is the fundamental requirement that distinguishes this specific complex from the others.

Crystal field theory models ligands as simple point charges, a drastic oversimplification that fails to account for the specific electronic interactions that dictate ligand strength, thereby leaving the spectrochemical series empirically observed rather than fundamentally explained. Furthermore, valence bond theory is useful for predicting geometry and magnetic behavior, but it provides no quantitative measure for the kinetic stability of coordination compounds. The nickel complex $[\mathrm{Ni}(\mathrm{CN})_4]^{2-}$ involves a $d^8$ metal center where the strong cyanide ligands force electron pairing to achieve $dsp^2$ hybridization, resulting in a square planar geometry. Finally, the cis-isomer of $[\mathrm{PtCl}_2(\mathrm{en})_2]^{2+}$ is optically active and exists as a pair of enantiomers, meaning the claim that there is only one such isomer is incorrect.

Analytical chemistry often utilizes specific coordination reactions to detect transition metals, where organic ligands create stable, insoluble complexes with distinct color profiles. Dimethyl glyoxime acts as a selective chelating agent that reacts with nickel ions in a basic medium to form a complex through the equation $\mathrm{Ni}^{2+}+2\mathrm{dmg}^{-}\longrightarrow [\mathrm{Ni}(\mathrm{dmg})_2]$. This resulting product is a characteristic rosy or brilliant red precipitate, providing a reliable and visually striking diagnostic test for identifying the presence of nickel in a chemical sample.

Many chemical substances and biological compounds owe their specific reactivity or biological function to the central metal atom they coordinate. In industrial processes, the Ziegler catalyst utilizes Titanium to facilitate polymer synthesis, while the Wilkinson catalyst relies on the specific coordination geometry of Rhodium to drive hydrogenation reactions. Biological systems also employ metal centers for critical tasks, such as Iron serving as the oxygen-carrying site in blood pigments and Cobalt acting as the essential metal ion within the complex structure of Vitamin $\mathrm{B}_{12}$.

The color of transition metal complexes often arises from d-d transitions, which require the presence of unpaired electrons within the d-orbitals. In the octahedral $[\mathrm{Ni}(\mathrm{H}_2\mathrm{O})_6]^{2+}$ complex, the nickel ion retains unpaired electrons, allowing these transitions to occur and giving the solution its characteristic green color. In contrast, the $[\mathrm{Ni}(\mathrm{CN})_4]^{2-}$ complex features strong-field ligands that force the pairing of all d-electrons, resulting in a diamagnetic square planar geometry where no d-d transitions are possible, thus leaving the solution colorless.

The reaction sequence begins with a group-IV cation that forms a black precipitate upon interaction with $\mathrm{H_2S}$, identifying the species as cobalt(II) sulfide, $\mathrm{CoS}$. Dissolving this precipitate in aqua regia yields $\mathrm{Co}^{2+}(\mathrm{aq})$ ions, which then undergo a transformation in the presence of $\mathrm{KNO}_2$ and $\mathrm{CH_3COOH}$ to form the stable complex $\mathrm{K}_3[\mathrm{Co}(\mathrm{NO}_2)_6]$. Identifying the central metal ion as $\mathrm{Co}^{3+}$ is essential here, as the oxidation state increases during the formation of this complex.

In this coordination environment, cobalt(III) possesses a $3d^6$ electronic configuration. The nitrite ions act as strong field ligands, which induce electronic pairing within the metal's $d$-orbitals to form a low-spin octahedral complex with $d^2sp^3$ hybridization. Because all six $d$-electrons are paired within the lower energy orbitals, the number of unpaired electrons is zero, leading to a spin-only magnetic moment calculated as $\sqrt{0(0+2)}$, which equals 0 BM.

Lead chromate exhibits amphoteric properties, allowing it to dissolve in concentrated, hot alkali solutions to form a soluble plumbite complex. The chemical reaction between lead chromate and excess hot $\mathrm{NaOH}$ yields the tetrahydroxoplumbate(II) ion, expressed as $[\mathrm{Pb}(\mathrm{OH})_4]^{2-}$. This specific species features a coordination number of four because four hydroxide ligands coordinate around the central lead atom, resulting in a dianionic complex.