Coordination Compounds – NEET UG ChemistryPractice Questions & PYQs

The wavelength of light absorbed during electronic transitions is inversely proportional to the crystal field splitting energy of the coordination complex, meaning stronger ligands and higher oxidation states lead to larger splitting parameters and shorter absorption wavelengths. Evaluating the series requires examining the ligand field strength and metal ion characteristics for each entity. The cyanide complex $[\mathrm{Co}(\mathrm{CN})_6]^{3-}$ contains a very strong-field ligand, resulting in a large splitting energy and the shortest absorption wavelength at 310 nm. Ammonia in $[\mathrm{Co}(\mathrm{NH}_3)_6]^{3+}$ provides intermediate field strength, corresponding to an absorption wavelength of 475 nm. The aquated titanium complex $[\mathrm{Ti}(\mathrm{H}_2\mathrm{O})_6]^{3+}$ absorbs at 498 nm due to a slightly smaller splitting parameter. Finally, $[\mathrm{Cu}(\mathrm{H}_2\mathrm{O})_4]^{2-}$ exhibits the smallest crystal field splitting among these species, yielding the longest absorption wavelength at 600 nm. Arranging these observed values in increasing order establishes the correct sequence from shortest to longest wavelength.

Ligands are classified by their denticity, which corresponds to the number of donor atoms available to coordinate simultaneously with a central metal ion. Ethylene diaminetetraacetate contains two nitrogen atoms and four oxygen atoms capable of donating lone pairs, providing a total of six binding sites. Because it attaches to the metal center through all six of these donor atoms, it functions as a hexadentate ligand.

Coordination compounds exhibit various forms of structural isomerism depending on ligand connectivity and the distribution of species inside versus outside the coordination sphere. The complex $[\mathrm{Co}(\mathrm{NH}_3)_5(\mathrm{NO}_2)]\mathrm{Cl}_2$ features an ambidentate ligand capable of coordinating through different donor atoms, resulting in linkage isomerism. Meanwhile, $[\mathrm{Co}(\mathrm{NH}_3)_5(\mathrm{SO}_4)]\mathrm{Br}$ exchanges ions between the coordination sphere and counter ions to display ionization isomerism, and $[\mathrm{Co}(\mathrm{NH}_3)_6][\mathrm{Cr}(\mathrm{CN})_6]$ contains complex cations and anions that interchange ligands to produce coordination isomerism. Lastly, $[\mathrm{Co}(\mathrm{H}_2\mathrm{O})_6]\mathrm{Cl}_3$ involves water molecules alternating between coordination positions and the outer lattice, which corresponds to solvate isomerism.

Ambidentate ligands are species that possess multiple distinct donor atoms but coordinate to a metal center through alternative atoms depending on the reaction conditions. While species like $\mathrm{SCN}^{-}$, $\mathrm{NO}_2^{-}$, and $\mathrm{CN}^{-}$ display this versatile dual-bonding capability, the oxalate ion ($\mathrm{C}_2 \mathrm{O}_4^{2-}$) acts as a bidentate ligand by engaging both of its oxygen atoms simultaneously rather than choosing between different donor atoms. Consequently, it does not qualify as an ambidentate ligand.

Dinuclear metal carbonyl complexes are characterized by unique bonding arrangements where transition metals possess a formal oxidation state of zero due to the neutral nature of carbon monoxide ligands. In structures such as $[\mathrm{Mn}_2(\mathrm{CO})_{10}]$ and $[\mathrm{Co}_2(\mathrm{CO})_8]$, the metal centers achieve electronic stability by forming a direct metal-metal covalent bond alongside multiple terminal carbonyl groups. Because these specific frameworks do not require bridging carbonyl groups to satisfy their coordination requirements, only the metal-metal linkage, terminal ligands, and zero oxidation state attributes apply.

Crystal field splitting and the chemical nature of surrounding ligands determine whether an octahedral transition metal complex adopts a low-spin paired configuration or a high-spin unpaired arrangement. In both $\mathrm{[Co(NH_3)_6]^{3+}}$ and $\mathrm{[CoF_6]^{3-}}$, the cobalt ion is in the $+3$ oxidation state with a $d^6$ electronic configuration, yet their magnetic profiles differ entirely. Ammonia functions as a strong-field ligand that forces electron pairing to produce a diamagnetic complex, while fluoride acts as a weak-field ligand that leaves unpaired electrons, rendering the system paramagnetic.

Coordination complexes are categorized as homoleptic when the central metal ion is bound exclusively to a single type of donor group, and heteroleptic when coordinated to two or more different types of ligands. The ion $[\mathrm{Co}(\mathrm{NH}_3)_6]^{3+}$ features only ammonia molecules surrounding the metal center, satisfying the definition of a homoleptic species. In contrast, $[\mathrm{Co}(\mathrm{NH}_3)_4\mathrm{Cl}_2]^{+}$ incorporates both ammine and chloride ligands within its coordination sphere, making it a heteroleptic complex.

Double salts completely lose their ionic identity by dissociating into all constituent simple ions when dissolved in water, whereas coordination compounds retain their complex entities in solution. Potash alum, represented as $\mathrm{KAl}(\mathrm{SO}_4)_2\cdot12\mathrm{H}_2\mathrm{O}$, is a typical double salt that yields free aluminum, potassium, and sulfate ions. On the other hand, $\mathrm{K}_2[\mathrm{Pt}(\mathrm{NH}_3)_2\mathrm{Cl}_2]$ contains a coordination sphere where the platinum metal center is firmly bound to ligands that do not dissociate individually.

The chelate effect provides exceptional thermodynamic stability to coordination compounds containing multidentate ligands because forming ring structures with the metal ion increases entropy and decreases dissociation rates. Among the given choices, ethylenediamine acts as a chelating ligand that forms stable five-membered rings with the cobalt center. Consequently, the complex featuring these chelating rings exhibits a much higher stability than the complexes possessing only monodentate ligands.

Homoleptic complexes are characterized by a central metal atom bound exclusively to a single type of donor ligand, meaning no mixed ligands are present in the coordination sphere. Potassium trioxalatoaluminate (III), formulated as $\mathrm{K}_3[\mathrm{Al}(\mathrm{C}_2\mathrm{O}_4)_3]$, features an aluminum center coordinated solely by three bidentate oxalate ions. Because every donor atom attached to the metal belongs to this single oxalate ligand class, the compound qualifies as a homoleptic complex.

Solvate isomerism occurs when coordination compounds have the same empirical formula but differ in whether water molecules act as ligands directly bound to the metal center or as free molecules of crystallization outside the coordination sphere. In each of these chromium complexes, six ligands are directly attached to the chromium ion, resulting in a fixed coordination number of six regardless of water exchange. This shifting distribution of water molecules between the inside and outside of the coordination bracket defines the solvate isomerism shown by these species.

Boron has a small atomic size and lacks vacant $d$-orbitals in its valence shell, which strictly limits its maximum covalency to four. Because of this restriction, boron cannot expand its octet to coordinate six water molecules around itself to form a species analogous to $\left[\mathrm{M}(\mathrm{H}_2\mathrm{O})_6\right]^{3+}$. In contrast, heavier group 13 elements like gallium, indium, and aluminium readily utilize their available $d$-orbitals or larger size to achieve a coordination number of six.

The color of coordination complexes arises from $d-d$ electronic transitions, where the energy absorbed corresponds directly to the crystal field splitting parameter dictated by the surrounding ligands. According to the spectrochemical series, ethylenediamine creates a significantly stronger ligand field than water molecules, resulting in greater orbital splitting. Consequently, complexes possessing a higher number of ethylenediamine ligands demand higher energy absorption for electronic promotion. Comparing the ligand compositions across the three complexes establishes that $[\mathrm{Ni}(\mathrm{en})_{3}]^{2+}$ absorbs the highest energy, followed by $[\mathrm{Ni}(\mathrm{H}_{2}\mathrm{O})_{2}(\mathrm{en})_{2}]^{2+}$, with $[\mathrm{Ni}(\mathrm{H}_{2}\mathrm{O})_{4}(\mathrm{en})]^{2+}$ absorbing the least.

Naming coordination entities requires identifying the cationic and anionic parts, listing the ligands alphabetically, and specifying the metal oxidation state in Roman numerals. The cation $[\mathrm{Ag}(\mathrm{H}_2\mathrm{O})_2]^+$ contains two neutral aqua ligands attached to silver, whereas the anion $[\mathrm{Ag}(\mathrm{CN})_2]^-$ contains two cyanido ligands attached to an anionic silver center designated by the suffix "-ate". Calculating the charge on each silver atom confirms an oxidation state of $I$ in both components of the salt. Arranging these parts yields diaquasilver(I) for the cation and dicyanidoargentate(I) for the anion.

Isomerism in coordination compounds arises when complexes share identical molecular formulas but differ in structural bonding or spatial distribution. The presence of the ambidentate ligand in $\mathrm{[Co(NH_3)_5NO_2]Cl_2}$ and $\mathrm{[Co(NH_3)_5ONO]Cl_2}$ permits linkage isomerism through alternative donor atom attachment, while the interchange of ligands between cationic and anionic coordination spheres in $\mathrm{[Cr(NH_3)_6][Co(CN)_6]}$ and $\mathrm{[Cr(CN)_6][Co(NH_3)_6]}$ generates coordination isomerism. Exchanging sulfate and bromide ions between coordination and ionization spheres in $\mathrm{[Co(NH_3)_5(SO_4)]Br}$ and $\mathrm{[Co(NH_3)_5Br]SO_4}$ produces ionization isomerism. A difference in whether water molecules serve as ligands inside the sphere or as solvent molecules of crystallization distinguishes the solvate isomers $\mathrm{[Cr(H_2O)_6]Cl_3}$ and $\mathrm{[Cr(H_2O)_5Cl]Cl_2 \cdot H_2O}$.

Synergistic bonding in metal carbonyls imparts dual $\sigma$ and $\pi$ bonding characteristics to the metal-carbon interaction, making the assertion entirely accurate. This bonding mechanism relies on a coordinate covalent interaction where the carbon monoxide ligand donates electron density into a vacant metal orbital to form a $\sigma$ bond. Simultaneously, electron density is back-donated from filled metal $d$-orbitals into the empty $\pi^*$ antibonding orbitals of the carbonyl ligand, constituting a $\pi$ bond. Because the reason incorrectly reverses these directional roles by claiming the ligand-to-metal bond is a $\pi$ bond, the reason statement fails while the assertion remains correct.

The spin-only magnetic moment of a coordination complex is determined using the relation $\mu = \sqrt{n(n+2)}$, where $n$ denotes the number of unpaired electrons. For the hexacyanoferrate(III) ion, $\mathrm{Fe}(\mathrm{CN})_{6}^{3-}$, iron is in the $\mathrm{Fe}^{3+}$ oxidation state with a $d^5$ configuration, and the strong-field cyanide ligand forces pairing so that only $1$ unpaired electron remains, corresponding to a magnetic moment of $1.73\text{ BM}$.

In the case of hexaaquairon(III), $\mathrm{Fe}(\mathrm{H}_{2}\mathrm{O})_{6}^{3+}$, the metal ion is again $\mathrm{Fe}^{3+}$ ($d^5$), but water acts as a weak-field ligand that prevents pairing, leaving $5$ unpaired electrons and resulting in a high magnetic moment of $5.92\text{ BM}$.

For the hexacyanoferrate(II) ion, $\mathrm{Fe}(\mathrm{CN})_{6}^{4-}$, the $\mathrm{Fe}^{2+}$ ion has a $d^6$ configuration where the strong cyanide ligands cause complete pairing, yielding $0$ unpaired electrons and a magnetic moment of $0\text{ BM}$. Finally, hexaaquairon(II), $\mathrm{Fe}(\mathrm{H}_{2}\mathrm{O})_{6}^{2+}$, contains $\mathrm{Fe}^{2+}$ paired with weak-field water ligands that do not induce pairing, leaving $4$ unpaired electrons and generating a magnetic moment of $4.90\text{ BM}$.

Ligand denticity is determined by the total number of electron-donating atoms a single molecule or ion uses to bind simultaneously to a central metal ion. The ethylenediaminetetraacetate structure incorporates two nitrogen atoms from the central diamine backbone and four oxygen atoms from the negatively charged acetate arms. Each of these six atoms possesses a lone pair of electrons available for coordination, allowing the ion to wrap completely around a metal center. This unique architecture establishes EDTA as a classic hexadentate chelating agent.

The spectrochemical series arranges ligands in order of their increasing crystal field splitting ability, reflecting how strongly they interact with the central metal ion's d-orbitals. Among the given species, thiocyanate ($\mathrm{SCN}^{-}$) acts as a relatively weak field ligand, followed by fluoride ($\mathrm{F}^{-}$) and oxalate ($\mathrm{C}_2\mathrm{O}_4^{2-}$), while cyanide ($\mathrm{CN}^{-}$) sits further along the series as a strong field ligand. Tracing these positions provides the increasing field strength sequence of $\mathrm{SCN}^{-} < \mathrm{F}^{-} < \mathrm{C}_2\mathrm{O}_4^{2-} < \mathrm{CN}^{-}$.

The spin-only magnetic moment of a transition metal ion is determined by the formula $\mu = \sqrt{n(n+2)}$, where $n$ denotes the count of unpaired electrons. Neutral chromium has an electronic configuration of $[\mathrm{Ar}]3d^5 4s^1$, meaning the $\mathrm{Cr}^{2+}$ ion is formed by removing two electrons to yield a $[\mathrm{Ar}]3d^4$ configuration. Counting the electrons in the $d$-orbitals reveals $n = 4$ unpaired electrons available for bonding. Substituting this value into the equation yields $\mu = \sqrt{4(4+2)} = \sqrt{24}$, which evaluates to $4.90$ BM.