D And F Block Elements – JEE Main ChemistryPractice Questions & PYQs

To determine the highest oxidation state among the given options, we calculate the charge of each metal center by balancing it against the charges of the attached ligands or oxygen atoms. In $MnO_2$, the manganese atom balances two oxide ions, resulting in an oxidation state of +4, while in $CrO_2Cl_2$, the chromium atom balances two oxide ions and two chloride ions, reaching an oxidation state of +6. Because these calculated values for manganese and chromium are significantly higher than the oxidation states of the metals in the other listed complexes, such as the +2 and +3 states found in the various ionic and coordination compounds, this pair represents the highest oxidation states.

The oxidation-reduction behavior of zinc in nitric acid is fundamentally dependent on the concentration of the acid, as this determines the extent to which the nitrate ion is reduced. When zinc reacts with dilute nitric acid, the reduction leads to the formation of nitrous oxide gas, represented by the following chemical equation:
$4\mathrm{Zn} + 10\mathrm{HNO}_3 \rightarrow 4\mathrm{Zn}(\mathrm{NO}_3)_2 + \mathrm{N}_2\mathrm{O} + 5\mathrm{H}_2\mathrm{O}$
Conversely, in the presence of concentrated nitric acid, the reaction is more aggressive and drives the reduction to nitrogen dioxide, as shown here:
$\mathrm{Zn} + 4\mathrm{HNO}_3 \rightarrow \mathrm{Zn}(\mathrm{NO}_3)_2 + 2\mathrm{H}_2\mathrm{O} + 2\mathrm{NO}_2$
Consequently, the reaction of zinc with dilute and concentrated nitric acid, respectively, yields nitrous oxide and nitrogen dioxide as the primary nitrogenous products.

The phenomenon of color in many gemstones arises from d-d electronic transitions within transition metal ions trapped in a crystal lattice, where the surrounding crystal field splits the energy levels of the d-orbitals. In the case of ruby, which is a variety of corundum, and emerald, a type of beryl, the vibrant red and green colors are both caused by the presence of chromium ions in the +3 oxidation state. Therefore, both gemstones derive their unique coloration from the inclusion of $Cr^{3+}$ ions, which absorb specific visible light wavelengths determined by the host crystal structure.

Transition metal oxides display a wide range of electrical and magnetic behaviors depending on how their d-orbitals overlap and interact with oxygen. In the case of $CrO_2$, the specific overlap of these orbitals facilitates the delocalization of electrons, resulting in metallic conductivity, while the alignment of magnetic moments creates the property of ferromagnetism. In contrast, other common oxides like $VO_2$, $MnO_2$, and $TiO_2$ do not exhibit this specific combination of metallic character and ferromagnetic ordering under standard conditions.

The stability of an ion in an aqueous environment is primarily dictated by its inherent tendency to undergo disproportionation or react with the solvent. While $\mathrm{Cu}^{+}$ ions are unstable in water because they spontaneously disproportionate into $\mathrm{Cu}^{2+}$ and copper metal according to the reaction $2\mathrm{Cu}^{+}\rightarrow\mathrm{Cu}^{2+}+\mathrm{Cu}$, other species possess specific environmental requirements for stability. Although $\mathrm{MnO_4^{2-}}$ is susceptible to disproportionation under neutral or acidic conditions, it remains stable in strongly alkaline aqueous solutions, which allows it to exist as the most stable species among the options provided.

Industrial synthesis relies on specific catalytic systems to drive production efficiency, where $\mathrm{TiCl}_3$ is a key component in the Ziegler-Natta polymerization process. The Wacker process utilizes $\mathrm{PdCl}_2$ for chemical oxidation, while the Deacon process relies on $\mathrm{CuCl}_2$ as an essential agent for the industrial production of chlorine gas. Completing these fundamental chemical pathways, $\mathrm{V_2O_5}$ acts as the necessary catalyst in the contact process for the large-scale manufacturing of sulfuric acid.

In potassium permanganate ($\mathrm{KMnO}_4$), the central manganese atom exists in the $+7$ oxidation state with a $d^0$ electronic configuration, meaning there are no electrons available for typical $d-d$ transitions to occur. The intense purple color arises instead from the electronic excitation where electrons transfer from the filled orbitals of the oxygen ligands to the empty $d$-orbitals of the metal center. This specific process is defined as an L → M charge transfer transition, which effectively absorbs visible light and produces the characteristic deep color of the permanganate ion.

Predicting the feasibility of these reaction sequences relies on understanding the thermal stability of iron salts and the reduction potential of various agents. For instance, heating $\mathrm{FeCl}_3$ in the presence of air does not lead to the formation of $\mathrm{FeCl}_2$, making that specific process invalid. Similarly, attempting to isolate metallic iron by simply heating $\mathrm{Fe}_2(\mathrm{SO}_4)_3$ or $\mathrm{FeSO}_4$ fails because these salts thermally decompose into iron oxides rather than the desired metal. Conversely, the conversion pathway shown in B, where iron reacts with oxygen to form $\mathrm{Fe}_3\mathrm{O}_4$ and is subsequently reduced in a stepwise manner by carbon monoxide at 600 °C and 700 °C, accurately describes valid metallurgical reduction processes.

The standard electrode potential $E^0_{M^{3+}/M^{2+}}$ measures the thermodynamic stability of trivalent ions relative to their divalent state, which is primarily governed by the electronic configuration stability of the transition metal ions. Cobalt exhibits the highest reduction potential of $1.97\text{ V}$ because the $Co^{3+}$ ion, possessing a $3d^6$ configuration, is significantly less stable than the $Co^{2+}$ ion with its $3d^7$ configuration, thereby making the reduction process highly favorable. In contrast, manganese shows a high value of $1.57\text{ V}$ as it moves toward the exceptionally stable half-filled $d^5$ shell, while iron at $0.77\text{ V}$ and chromium at $-0.41\text{ V}$ display lower potentials due to the inherent stability of their respective $3+$ oxidation states.

Paramagnetic behavior in transition metal ions is directly proportional to the number of unpaired electrons in the $d$-subshell. Evaluating the ions in the sequence, $\mathrm{Mn}^{2+}$ has a $\mathrm{3d^5}$ configuration, which corresponds to 5 unpaired electrons. In contrast, $\mathrm{Fe}^{2+}$ has a $\mathrm{3d^6}$ configuration, which results in 4 unpaired electrons because the first five electrons occupy the $d$-orbitals singly before the sixth one pairs up. Since $\mathrm{Mn}^{2+}$ possesses more unpaired electrons than $\mathrm{Fe}^{2+}$, the proposed order claiming that $\mathrm{Fe}^{2+}$ is more paramagnetic is incorrect. The other trends provided are accurate: ionic radii generally decrease across the transition period due to increasing effective nuclear charge, manganese exhibits the most extensive range of oxidation states due to the stability of its half-filled $d$-subshell and available $4s$ electrons, and the stability of trivalent cations in aqueous solution, which relates to their hydration enthalpies, identifies $\mathrm{Sc}^{3+}$ as the least stable because it has the lowest hydration enthalpy ($\mathrm{Sc}^{3+} = 3960\,\text{kJ/mol}$) compared to $\mathrm{Cr}^{3+}$ ($\mathrm{Cr}^{3+} = 4563\,\text{kJ/mol}$), $\mathrm{Fe}^{3+}$ ($\mathrm{Fe}^{3+} = 4429\,\text{kJ/mol}$), and $\mathrm{Co}^{3+}$ ($\mathrm{Co}^{3+} = 4653\,\text{kJ/mol}$).

Hydrolysis of metal ions in aqueous solution is primarily governed by the charge density of the cation, which determines how strongly it polarizes coordinated water molecules. Since the ferric ion ($Fe^{3+}$) possesses a higher positive charge compared to the ferrous ion ($Fe^{2+}$), it exerts a stronger electrostatic attraction that pulls electron density away from the water ligands more effectively. This enhanced polarization promotes the dissociation of protons from the water molecules, meaning ferric compounds are significantly more prone to hydrolysis than ferrous compounds, rendering the assertion that ferrous compounds hydrolyze more easily incorrect.

Lanthanoids are fundamentally defined by the stability of the $+3$ oxidation state, which serves as the most prominent state across the entire series. While specific elements may display a $+4$ oxidation state under certain conditions, it is chemically inaccurate to suggest that this is a universal feature shared by every single member of the group. The similarity in chemical behavior resulting from this common $+3$ state makes isolating individual lanthanoids remarkably difficult, just as the gradual decrease in ionic radii across the series remains a consistent hallmark of their chemistry.

The stability of an atom is significantly increased by the symmetry and exchange energy associated with half-filled subshells, a principle that governs the electronic arrangement of Gadolinium. Because a half-filled $4f$ subshell provides extra stability, the electron configuration deviates from the standard Aufbau sequence as the 64th electron enters the $5d$ orbital instead of continuing to fill the $4f$ subshell. This unique arrangement results in the final outer electron configuration of $4f^7 5d^1 6s^2$.

Transition elements exhibit variable oxidation states, and the chemical nature of their oxides and hydroxides is directly governed by these specific states. Lower oxidation states typically result in basic character, whereas higher oxidation states increasingly exhibit acidic properties. Consequently, the claim that transition metals show basic character in their highest oxidation states contradicts this fundamental principle, as such high-energy states are instead associated with acidic behavior.

The trivalent oxidation state of lanthanoids is governed by their electronic configuration, particularly the occupancy of their $f$-orbitals. Because most $Ln^{3+}$ ions contain electrons in the $f$-subshell that facilitate $f$-$f$ electronic transitions, these compounds typically exhibit distinct colors rather than being colorless. While $La^{3+}$ and $Lu^{3+}$ are notable exceptions due to their stable $f^0$ and $f^{14}$ configurations respectively, asserting that lanthanoid compounds are generally colorless ignores the prevalent behavior of the majority of the series.

In a redox titration, the oxidizing agent must react exclusively with the analyte to ensure precise stoichiometric measurements. Potassium permanganate ($\mathrm{KMnO}_4$) is a potent oxidant that lacks the selectivity required to distinguish between the oxalic acid and the chloride ions present in hydrochloric acid ($\mathrm{HCl}$). This leads to a side reaction where the permanganate oxidizes the $\mathrm{HCl}$ into chlorine gas ($\mathrm{Cl}_2$) while being reduced itself to $\mathrm{Mn}^{2+}$, which ultimately consumes the titrant unnecessarily and invalidates the results.

The variety of oxidation states an element exhibits is determined by the ease with which electrons in various orbitals can participate in chemical bonding. In lanthanoids, the $4f$ orbitals are deeply buried, creating a large energy barrier between the $4f$ and $5d$ orbitals that limits the variability of their oxidation states. Actinoids, however, have $5f$ orbitals that are less shielded, leading to a significantly smaller energy difference between the $5f$ and $6d$ orbitals, which allows these electrons to more readily participate in bonding and thus generate a wider range of oxidation states.

The variety of oxidation states displayed by an element is governed by the strength of the electrostatic interaction between the nucleus and the outer orbitals. Because the $5f$ orbitals in actinoids extend further from the nucleus than the $4f$ orbitals in lanthanoids, the effective nuclear pull on these electrons is notably weaker. This reduced attraction makes the $5f$ electrons more available for chemical bonding, which allows actinoids to demonstrate a broader spectrum of oxidation states than lanthanoids.

The effectiveness of electron shielding depends heavily on the spatial proximity of an orbital to the nucleus. Because the $4f$ orbital is situated closer to the nucleus than the $5f$ orbital, it provides a superior screening effect for the valence electrons. Consequently, the assertion that $4f$ and $5f$ orbitals possess equal shielding capabilities is incorrect, as their distinct radial distances result in significantly different screening abilities.

Nickel possesses an atomic number of 28, corresponding to an electronic configuration of $[Ar] 3d^8 4s^2$ for the neutral atom. Forming the $\mathrm{Ni}^{2+}$ ion involves losing the two $4s$ electrons, which results in a $3d^8$ configuration. With water acting as a weak field ligand, no pairing of the electrons occurs in the $3d$ orbitals, leaving two unpaired electrons. The spin-only magnetic moment follows the relationship $\mu = \sqrt{n(n+2)} \mu_B$, where $n$ is the number of unpaired electrons, and substituting $n = 2$ results in $\mu = \sqrt{2(2+2)} = \sqrt{8}$, which is approximately 2.84 $\mu_B$.