Coordination Compounds – NEET UG ChemistryPractice Questions & PYQs

Chelating ligands form stable coordinate ring structures with metal ions, with ring size and stability heavily influenced by the ligand backbone geometry. Acetylacetonate coordinates through two oxygen atoms to the central cobalt ion, creating a stable six-membered ring that encompasses the metal center and the carbon-oxygen framework of the ligand. This specific ring size minimizes angular strain and maximizes stability through conjugated double bond resonance.

The spin-only magnetic moment of a transition metal complex is determined by the formula $\mu = \sqrt{n(n+2)}$ BM, where $n$ represents the number of unpaired electrons. In the complex $\mathrm{Cu}(\mathrm{NH}_3)_4^{2+}$, copper exhibits an oxidation state of $+2$, which corresponds to a $3d^9$ electronic configuration with exactly one unpaired electron. Substituting this single unpaired electron into the magnetic moment expression yields $\sqrt{1(1+2)}$, which evaluates to $\sqrt{3}$ or 1.73 BM.

Tetrahedral coordination entities invariably adopt a high-spin configuration because the crystal field splitting energy is considerably smaller than the pairing energy, leaving orbitals singly occupied. In the complex $\mathrm{[CoBr_4]^{2-}}$, cobalt exists in the $+2$ oxidation state, corresponding to a $3d^5$ electronic configuration as noted in the reference parameters. Due to the weak-field nature of bromide ligands combined with the inherently small tetrahedral splitting, the electrons remain unpaired across the $d$-orbitals, imparting a distinct paramagnetic character to the species.

Crystal field stabilization energy for octahedral complexes is calculated using the relation $\text{CFSE} = (-0.4 x + 0.6 y) \Delta_0$, where $x$ and $y$ denote electron counts in the $t_{2g}$ and $e_g$ sets. For a high-spin $d^4$ configuration, pairing energy exceeds crystal field splitting, causing the fourth electron to occupy the higher energy $e_g$ orbital rather than pair up. Substituting the resulting electron distribution of three electrons in the lower set and one in the upper set into the energy expression results in a value of $-0.6 \Delta_0$.

Dissolving a coordination compound releases its outer-sphere counter-ions into solution where they can undergo precipitation reactions with reagents like silver nitrate. The compound dichlorotetraaquachromium (III) chloride ionizes to release one chloride ion per formula entity into the solution. Multiplying the given molarity of 0.01 M by the volume of 100 mL yields 0.001 moles of the coordination complex, which directly dictates that 0.001 moles of silver chloride precipitate upon adding excess silver nitrate.

Strong-field ligands possess a high crystal field splitting energy that forces electrons to pair up within lower energy orbitals rather than remaining singly distributed. For the square planar entity $\mathrm{[Ni(CN)_4]^{2-}}$, nickel exhibits a $+2$ oxidation state with a $3d^8$ electronic configuration. The strong cyanide ligands cause complete pairing of these $d$-electrons, resulting in fully occupied orbitals with zero unpaired electrons and a diamagnetic response.

Coordination nomenclature requires naming anionic ligands alphabetically, followed by neutral ligands using multiplicative prefixes like bis, then the central metal with its Roman numeral oxidation state, and finally the counter anion. The complex cation contains two fluoride ligands designated as difluorido and two ethylenediamine ligands designated using the prefix bis, while the single chloride counter-ion balances the coordination sphere to establish a $+3$ oxidation state for chromium. Assembling these components gives the complete IUPAC designation of difluoridobis(ethylenediamine)chromium(III) chloride.

Crystal field stabilization energy for octahedral complexes is determined by populating the lower energy $t_{2g}$ and higher energy $e_g$ orbitals, where each electron in the $t_{2g}$ set contributes $-\frac{2}{5}\Delta_o$. For a low-spin $d^6$ configuration, all six electrons enter and completely fill the $t_{2g}$ level, forming three electron pairs in the process. Multiplying the number of electrons by their orbital energy and adding the pairing energy for the three formed pairs yields the total expression $\frac{-12}{5}\Delta_o + 3P$.

Outer orbital coordination complexes utilize their higher energy $d$-orbitals for hybridization, typically forming $sp^3d^2$ manifolds when bonding with ligands in an octahedral geometry. In the species $\mathrm{[Ni(NH_3)_6]^{2+}}$, the central metal ion accommodates six ammonia molecules by employing outer $4d$ orbitals alongside $4s$ and $4p$ orbitals, leaving the inner $3d$ electrons unforced to pair. Because the $3d$ shell retains unpaired electrons under this ligand field influence, the complex exhibits a pronounced paramagnetic property.

The reaction between ammoniacal nickel(II) and dimethylglyoxime produces the coordination complex $[\text{Ni(dmg)}_2]$, which adopts a square planar geometry rather than a tetrahedral one due to the $d^8$ electronic configuration of nickel in this environment. Dimethylglyoxime acts as a bidentate ligand, and the resulting chelate ring is further stabilized by strong intramolecular symmetrical hydrogen bonding. Consequently, the assertion that the red complex has a tetrahedral geometry is false.

Square planar complexes with four entirely different unidentate ligands, represented by the general formula $[Mabcd]$, always yield three distinct geometrical isomers due to different relative spatial arrangements. By fixing the position of one ligand and systematically rotating the remaining three positions around the central platinum atom, three unique configurations are generated. For the specific complex $[Pt(Py)(NH_3)BrCl]$, these correspond to pyridine being trans to ammonia, trans to bromine, and trans to chlorine. Thus, the restricted planar geometry allows for exactly three stereoisomers.

Coordination isomerism occurs in compounds containing both complex cations and complex anions when ligands are interchanged between the two metal centers. This structural rearrangement produces distinct chemical species even though the empirical formula remains entirely identical. The given pair features an exchange of ammine and cyanide ligands between the cobalt and chromium coordination spheres. Consequently, this mutual swapping of ligands classifies them precisely as coordination isomers.

Paramagnetic behavior is directly proportional to the number of unpaired electrons, denoted as $n$. Water acts as a weak field ligand in these octahedral complexes, allowing the $d$-electrons to populate orbitals while maximizing parallel spins according to Hund's rule. Analyzing the central metal ions, $Cr^{2+}$ has four unpaired electrons, $Mn^{2+}$ has five, $Fe^{2+}$ has four, and $Co^{2+}$ retains three unpaired electrons. Because the cobalt complex contains the lowest value of $n$, it exhibits the least degree of paramagnetism.

The strength of the carbon-oxygen bond in metal carbonyls is inversely related to the extent of $\pi$-backdonation from the metal's filled $d$-orbitals into the antibonding $\pi^*$ orbitals of the carbonyl ligand. When a metal complex carries a positive charge, its electron density decreases, which severely inhibits this ligand-directed back-donation. In the cationic species $\mathrm{Mn(CO)_6^+}$, the positive charge minimizes electron transfer to the $\mathrm{CO}$ ligands, preserving a higher bond order. Consequently, this lack of electronic crowding into antibonding orbitals produces the strongest $\mathrm{C-O}$ bond among the choices.

Diamagnetism in transition metal complexes arises when all constituent electrons are fully paired up within their respective orbitals, leaving no net magnetic moment. In the square planar tetracyanonickelate ion, nickel is present in a $+2$ oxidation state with a $d^8$ electronic configuration interacting with strong-field cyanide ligands. These strong ligands enforce complete pairing of the metal $d$-electrons, resulting in zero unpaired electrons and a purely diamagnetic behavior.

Paramagnetic behavior in coordination complexes is directly determined by the number of unpaired d-electrons present in the central metal ion. Titanium in $[\mathrm{Ti}(\mathrm{NH}_3)_6]^{3+}$ has a single unpaired electron, whereas chromium in $[\mathrm{Cr}(\mathrm{NH}_3)_6]^{3+}$ possesses three unpaired electrons in its subshell. In contrast, cobalt in $[\mathrm{Co}(\mathrm{NH}_3)_6]^{3+}$ and zinc in $[\mathrm{Zn}(\mathrm{NH}_3)_6]^{2+}$ contain no unpaired electrons due to complete electron pairing or a fully filled subshell. Consequently, $[\mathrm{Cr}(\mathrm{NH}_3)_6]^{3+}$ exhibits the highest paramagnetic behaviour because it maximizes the count of unpaired electrons.

Crystal field stabilization energy in octahedral coordination entities is calculated using the formula $\text{CFSE} = [-0.4n_{t_{2g}} + 0.6n_{e_g}]\Delta_o$, where $n_{t_{2g}}$ and $n_{e_g}$ represent the number of electrons in the respective $d$-subshell orbitals. For a high-spin $d^4$ configuration, three electrons occupy the lower-energy $t_{2g}$ set and one electron enters the higher-energy $e_g$ set. Substituting these electron counts into the expression gives $\text{CFSE} = 3(-0.4\Delta_o) + 1(0.6\Delta_o)$, which simplifies to $-1.2\Delta_o + 0.6\Delta_o$. Evaluating this summation yields a final crystal field stabilization energy of $-0.6\Delta_o$.

Tetrahedral coordination compounds of the type $[MA_2B_2]$ possess a high degree of structural symmetry where all ligand positions bear equivalent geometric relationships to one another. While square planar and octahedral geometries allow for distinct cis and trans spatial arrangements, tetrahedral complexes cannot support such alternate configurations. Consequently, the relative positions of the ligands in $[Ni(NH_3)_2Cl_2]$ cannot be altered to form distinct isomers, rendering it incapable of exhibiting isomerism.

Octahedral coordination entities of the formula $[MA_4B_2]$ can exist in cis and trans configurations, representing a spatial arrangement difference known as geometrical isomerism. When two identical ligands occupy adjacent positions, the isomer is termed cis, whereas opposite positions yield the trans isomer. These distinct spatial arrangements alter the absorption of light, directly accounting for the existence of two different coloured complexes for $[Co(NH_3)_4Cl_2]$.

Absorption of visible light by transition metal complexes relies on electronic transitions between d-orbitals, which require the presence of unpaired d-electrons to occur. Coordination complexes lacking unpaired electrons cannot facilitate these electronic transitions and therefore remain transparent to visible light. In the case of $[Fe(H_2O)_6]$, specific crystal field interactions and electron pairing yield a configuration devoid of unpaired electrons capable of absorbing visible wavelengths.