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

The actinoid series encompasses elements with atomic numbers from 89 to 103, characterized by the progressive filling of the $5f$ subshell. Since atomic number 101 falls squarely within this range, it is definitively categorized as an actinoid.

Following the completion of the actinoid series at atomic number 103, the subsequent elements represent the beginning of the transactinoid series. Atomic number 104 initiates this sequence by occupying the $6d$ subshell, positioning it in Group 4 of the periodic table, directly below zirconium and hafnium.

Mischmetal functions as a commercial alloy distinguished by its high concentration of lanthanoid metals, which constitute approximately 95% of its total composition. The remaining 5% is largely comprised of iron, alongside smaller trace amounts of sulfur, carbon, calcium, and aluminum. Because the vast majority of the material is composed of lanthanoids, it is fundamentally categorized as a lanthanoid-based alloy.

The spin-only magnetic moment of a transition metal ion is determined by its number of unpaired electrons $n$, according to the relationship $\mu = \sqrt{n(n+2)}$ B.M., where a higher value of $n$ results in a greater magnetic moment. Evaluating the electronic states, $Sc^{3+}$ has a $3d^0$ configuration with 0 unpaired electrons, $Ti^{3+}$ has a $3d^1$ configuration with 1 unpaired electron, $Ti^{2+}$ has a $3d^2$ configuration with 2 unpaired electrons, and $V^{2+}$ has a $3d^3$ configuration with 3 unpaired electrons. Following this relationship, the magnetic moments increase in the order $Sc^{3+} < Ti^{3+} < Ti^{2+} < V^{2+}$.

Interstitial compounds are formed when small atoms like hydrogen, carbon, or nitrogen become trapped within the voids of a transition metal crystal lattice, which significantly enhances the physical properties of the host metal. This structure grants the materials exceptional hardness, high melting points, and electrical conductivity similar to pure metals. Since these small atoms occupy the interstices without disrupting the fundamental metallic bonds, the resulting compounds are characterized by their chemical inertness rather than reactivity.

Transition elements generally exhibit variable oxidation states because the energy gap between the $(n-1)d$ and $ns$ orbitals is minimal, permitting electrons from both to participate in chemical bonding. While elements like Vanadium, Titanium, and Copper take advantage of their available d-electrons to exist in states such as $V(+2, +3, +4, +5)$, $Ti(+2, +3, +4)$, and $Cu(+1, +2)$, Scandium remains distinct. With an atomic configuration of $[Ar] 3d^1 4s^2$, Scandium loses these three electrons to attain the highly stable noble gas configuration of Argon, as removing further electrons from the inner shell would require excessive energy. Consequently, Scandium is restricted to the $+3$ oxidation state, unlike the other transition metals that readily adopt multiple configurations.

The phenomenon of lanthanoid contraction is the primary driver behind the nearly identical atomic radii observed in molybdenum and tungsten. Although atomic size typically increases when moving down a group, the filling of the $4f$ subshell between these two elements introduces poor shielding that causes the effective nuclear charge to rise significantly. This increased attraction on the valence electrons counteracts the expected growth in atomic size, resulting in the comparable radii found in this specific pair.

The Wilkinson catalyst is a quintessential organometallic complex utilized extensively in homogeneous catalysis for the selective hydrogenation of alkenes. Its structural identity is defined by a central rhodium metal atom coordinated to three triphenylphosphine ligands and one chloride atom. This specific chemical configuration is represented by the formula $[(Ph_3P)_3RhCl]$, which distinguishes the catalyst from analogous complexes involving different transition metals or alkyl-substituted phosphine ligands.

Ionization enthalpy generally increases across a transition metal period as the effective nuclear charge rises, resulting in a stronger electrostatic attraction that makes it more difficult to remove an electron. Based on experimental energy requirements, Titanium is the easiest to ionize at 656 kJ/mol, followed by Manganese at 717 kJ/mol, Nickel at 736 kJ/mol, and finally Zinc at 906 kJ/mol, which sits highest due to the extra stability of its fully filled $3d^{10}$ electronic configuration. This upward trend clearly illustrates the increasing energy needed to overcome nuclear attraction as one moves from Titanium toward Zinc.

The chemistry of manganese is defined by its versatile oxidation states, which are highlighted here through a series of redox transformations starting with the oxidation of manganese dioxide. When $\mathrm{MnO}_2$ (A) is fused with $\mathrm{KOH}$ in the presence of oxygen, it forms the green potassium manganate, $\mathrm{K}_2\mathrm{MnO}_4$ (B). This green compound subsequently disproportionates in an acidic medium, such as in the presence of $\mathrm{HCl}$, to produce purple potassium permanganate, $\mathrm{KMnO}_4$ (C), and regenerate $\mathrm{MnO}_2$.

The final stage of the sequence involves the strong oxidizing capability of $\mathrm{KMnO}_4$ (C) in an aqueous environment reacting with iodide ions. This reaction oxidizes iodide to iodate, forming potassium iodate, $\mathrm{KIO}_3$ (D), while simultaneously reducing the permanganate back to $\mathrm{MnO}_2$ (A). Mapping these chemical transformations confirms that the substance identified as A is $\mathrm{MnO}_2$ and the substance identified as D is $\mathrm{KIO}_3$.

Actinide elements demonstrate variable oxidation states because their 5f, 6d, and 7s orbitals are energetically close, allowing multiple electrons to participate in bonding. Uranium is characterized by a maximum oxidation state of 6, whereas plutonium can reach a higher oxidation state of 7 due to the electronic configuration of the later actinides. Consequently, the highest possible oxidation states for uranium and plutonium are 6 and 7, respectively.

The lanthanoid contraction arises because 4f orbitals possess a diffuse shape that provides ineffective shielding against the rising nuclear charge as the atomic number increases. This poor screening allows the positively charged nucleus to pull the electron cloud closer, resulting in a consistent reduction in size. Because of this persistent inward attraction, both the atomic radii and the ionic radii exhibit a steady, gradual decrease throughout the series.

Atomic radius is primarily dictated by the number of principal energy levels and the effective nuclear charge pulling electrons inward. Nitrogen, being a second-period element with fewer occupied shells, is distinctly smaller than the lanthanide elements. Within the lanthanides, Europium possesses the largest radius due to the electronic repulsions associated with its stable $4f^7$ configuration, whereas Cerium is larger than Holmium because the latter experiences the cumulative effects of lanthanide contraction as atomic number increases.

According to the Aufbau principle, electrons fill subshells in order of increasing $(n+l)$ energy, and because the $4f$ orbital ($n=4, l=3$) has a lower principal quantum number than the $5d$ orbital ($n=5, l=2$), it is prioritized during the filling process of the lanthanide series. The $4f$ subshell accommodates 14 electrons, and this filling progression spans from atomic number 58 to 71. Therefore, the 71st electron is the final electron required to complete the $4f$ subshell.

The enthalpy of atomization is primarily governed by the strength of metallic bonding, which depends on the number of unpaired electrons available for delocalization within the crystal lattice. While Zinc exhibits weak bonding, it is excluded from consideration here because its fully filled $d^{10}$ electronic configuration classifies it as a group 12 metal rather than a true transition element. Among the actual transition metals provided, Copper features the configuration $3d^{10} 4s^1$, which provides fewer unpaired electrons to contribute to bond strength compared to elements like Vanadium or Iron, resulting in the lowest enthalpy of atomization in the group.

The wide range of oxidation states exhibited by actinide elements is primarily driven by the small energy difference between their $5f$, $6d$, and $7s$ subshells, which allows for the involvement of multiple electrons in chemical bonding. Among the actinides, Neptunium and Plutonium demonstrate the maximum number of possible oxidation states because their specific electronic configurations allow for the most extensive participation of these orbitals compared to other elements in the series.

Ruby is formed by chromium ions substituting for aluminum ions in the alumina ($Al_2O_3$) crystal lattice, whereas the beryl mineral structure is entirely distinct and hosts chromium to form emeralds. Because ruby is an aluminum oxide rather than a silicate-based beryl structure, attributing the chromium sites of ruby to a beryl matrix is factually incorrect.

The remaining statements are accurate: $[\mathrm{Fe}(\mathrm{H}_2\mathrm{O})_6]^{2+}$ and $[\mathrm{Cr}(\mathrm{H}_2\mathrm{O})_6]^{2+}$ each contain four unpaired electrons resulting in similar magnetic moments, and $[\mathrm{CoCl}(\mathrm{NH}_3)_5]^{2+}$ appears violet as it absorbs yellow light. Furthermore, $[\mathrm{Ni}(\mathrm{NH}_3)_4(\mathrm{H}_2\mathrm{O})_2]^{2+}$ is an outer orbital complex with two unpaired electrons, yielding a calculated spin-only magnetic moment of $\sqrt{8} \approx 2.83$ BM.

The chemical process describes the characteristic oxidation of manganese dioxide ($MnO_2$) during alkali fusion with an oxidizing agent like potassium nitrate ($KNO_3$), which yields the dark green potassium manganate ($K_2MnO_4$). In an acidic environment, the green manganate ion ($MnO_4^{2-}$) is unstable and undergoes disproportionation into the deep purple permanganate ion ($MnO_4^-$) and manganese dioxide, uniquely identifying the transition metal $X$ as manganese.

The selective precipitation of metal sulfides in qualitative analysis relies on the solubility product ($K_{sp}$) of the sulfide and the concentration of $S^{2-}$ ions, which is severely suppressed in acidic media due to the common ion effect of $H^+$ from HCl. Copper ions ($Cu^{2+}$) are Group II cations that readily precipitate as black $CuS$ in these acidic conditions, whereas nickel ions ($Ni^{2+}$) are Group IV cations with a significantly higher $K_{sp}$ for $NiS$. Because $Ni^{2+}$ requires the higher sulfide ion concentration present in alkaline conditions to exceed its solubility product, it does not precipitate with $H_2S$ in the presence of HCl, making the statement that both ions form a black precipitate under these conditions incorrect.

The liberation of hydrogen gas by metal ions in an acidic medium depends on the standard reduction potential of the $M^{3+}/M^{2+}$ couple relative to the standard hydrogen electrode potential. For a metal ion to act as a reducing agent and liberate hydrogen via the reaction $2M^{2+} + 2H^+ \to 2M^{3+} + H_2$, the potential of the oxidation half-reaction $M^{2+} \to M^{3+} + e^-$ must be favorable, which occurs only if the standard reduction potential $E^\circ_{M^{3+}/M^{2+}}$ is negative. Since $Ti^{2+}$, $V^{2+}$, and $Cr^{2+}$ have negative standard reduction potentials, they are capable of reducing hydrogen ions, whereas $Mn^{2+}$ has a highly positive standard reduction potential ($E^\circ = +1.57$ V), rendering it stable and incapable of releasing hydrogen gas in the presence of dilute acids.

Zinc oxide exhibits amphoteric character, which allows it to react as an acid when paired with a basic oxide and as a base when paired with an acidic oxide. In the reaction $\mathrm{ZnO} + \mathrm{Na_2O} \rightarrow \mathrm{Na_2ZnO_2}$, zinc oxide acts as an acid because it reacts with the basic oxide $\mathrm{Na_2O}$ to form a salt. Conversely, in the reaction $\mathrm{ZnO} + \mathrm{CO_2} \rightarrow \mathrm{ZnCO_3}$, zinc oxide functions as a base by reacting with the acidic oxide $\mathrm{CO_2}$ to yield zinc carbonate.