P Block Elements – JEE Main ChemistryPractice Questions & PYQs

The structure of Buckminsterfullerene, denoted as $C_{60}$, is a cage-like molecule resembling a soccer ball that forms a truncated icosahedron, which contains exactly 12 pentagonal faces. In the case of white phosphorus, which exists as discrete $P_4$ molecules, the four phosphorus atoms are positioned at the corners of a regular tetrahedron, creating 4 triangular faces where each atom is bonded to the other three.

When iodine reacts with concentrated nitric acid, it undergoes oxidation to form iodic acid as the primary iodine-containing product. This chemical transformation is represented by the equation $I_2 + 10HNO_3 \rightarrow 2HIO_3 + 10NO_2 + 4H_2O$, where Y corresponds to $HIO_3$. By assigning the standard oxidation numbers of $+1$ for hydrogen and $-2$ for each oxygen atom, the oxidation state $x$ of iodine in this compound is determined by the equation $1 + x + 3(-2) = 0$, which results in a value of $+5$.

$\mathrm{B_{2}H_{6}}$ serves as a potent electron-deficient hydride that reacts aggressively with oxidizing agents and water due to the high thermodynamic stability of its oxidation products. Combustion in an oxygen-rich environment results in the oxidation of the boron centers to form boron trioxide along with water, expressed by the reaction
$\mathrm{B_{2}H_{6}} + 3\mathrm{O_{2}} \rightarrow \mathrm{B_{2}O_{3}} + 3\mathrm{H_{2}O}$
When it encounters liquid water, the molecule undergoes complete hydrolysis to produce boric acid and release hydrogen gas, following the stoichiometric equation
$\mathrm{B_{2}H_{6}} + 6\mathrm{H_{2}O} \rightarrow 2\mathrm{H_{3}BO_{3}} + 6\mathrm{H_{2}}$

Catenation refers to the ability of an element to form stable, self-linking covalent bonds, a property largely determined by the bond dissociation energy between atoms of the same element. Because bond strength is inversely related to atomic size, carbon exhibits the strongest catenation, while the capacity for chain formation decreases down the group as atomic radii increase and orbital overlap becomes less effective. This leads to the observation that $\mathrm{C} > \mathrm{Si} > \mathrm{Ge} \approx \mathrm{Sn}$, reflecting the systematic decline in bond stability among these elements.

The chemical nature of Group 13 oxides transitions from acidic to basic as metallic character increases down the group. Because boron is non-metallic, its trioxide, $B_2O_3$, is acidic, whereas aluminium and gallium occupy the intermediate range, forming amphoteric oxides like $Al_2O_3$ and $Ga_2O_3$ that interact with both acids and bases. Further down the group, indium and thallium exhibit distinct metallic properties, resulting in basic oxides, confirming that all three provided statements are correct.

The reactivity of halogens towards hydrogen decreases down the group, meaning iodine is significantly less reactive than fluorine, chlorine, and bromine. Consequently, the reaction $H_2 + I_2 \rightleftharpoons 2HI$ is a slow, endothermic process with a high activation energy that requires a catalyst, such as finely divided platinum, to proceed at a significant rate.

Noble gases such as helium, neon, and krypton are stable elements present in the Earth's atmosphere as trace components. In contrast, radon is a radioactive noble gas that is not available in the air, primarily because it is produced by the radioactive decay of other elements and undergoes further decay with a short half-life. Consequently, the option corresponding to this radioactive element is the correct choice for the noble gas absent from the atmosphere.

Silica ($SiO_2$) naturally exists in various forms, classified primarily by whether their internal atomic structure is ordered or disordered. While quartz, tridymite, and cristobalite form rigid, crystalline lattices with long-range repeating patterns, kieselguhr constitutes the amorphous variety of silica. Because it lacks a specific crystalline arrangement, kieselguhr stands out as the non-crystalline member among these options.

Catenation is fundamentally determined by the strength of the bonds an element forms with itself, a property that diminishes as atomic size increases. Within Group 14, the element-element bond dissociation energy decreases significantly as one progresses down the group, meaning larger atoms struggle to maintain stable, lengthy chains. Because Lead possesses the largest atomic radius among the elements listed, its bond strength is insufficient to support the formation of catenated structures, making it the only one among Silicon, Germanium, Tin, and Lead that does not exhibit this behavior.

The capability to form $p\pi - p\pi$ multiple bonds relies on the effectiveness of lateral overlap between atomic orbitals, which is significantly enhanced by a smaller atomic radius. Carbon is the only element in this group with a sufficiently small size to allow its 2p orbitals to achieve the tight, efficient sideways overlap required to form the strongest multiple bonds. As the atomic size increases down the group, the higher-energy orbitals become too diffuse to sustain stable $\pi$ bonding, rendering the process much less favorable for the heavier elements.

The increasing stability of the +1 oxidation state as we descend Group 13 is fundamentally driven by the inert pair effect, where the $ns^2$ electrons become increasingly reluctant to participate in chemical bonding due to poor shielding by intervening $d$ and $f$ orbitals. As the atomic size and shielding complexity grow, the energy penalty for utilizing these $s$-electrons in a +3 oxidation state rises, making the lower +1 oxidation state progressively more favored. This trend results in the stability of the +1 state increasing down the group in the sequence Al < Ga < In < Tl.

Buckminsterfullerene, or $\mathrm{C}_{60}$, consists of 60 carbon atoms arranged in a truncated icosahedral geometry to form a closed spherical cage. The distribution of its rings is governed by the geometric relationship $\frac{n}{2}-10$ for hexagons, where $n$ is the total number of carbon atoms, resulting in $\frac{60}{2}-10 = 20$ hexagonal rings alongside the 12 pentagonal rings required to maintain the spherical curvature.

Hydrides are classified as electron-deficient or electron-precise based on whether they possess enough valence electrons to satisfy bonding requirements through standard two-center, two-electron bonds. Group 13 hydrides such as $B_2H_6$, $AlH_3$, and $GaH_3$ lack the necessary valence electrons to complete these bonds without creating a deficiency. In contrast, silicon in $SiH_4$ provides exactly the right number of valence electrons to form four stable covalent bonds with hydrogen, making it an electron-precise hydride.

Piezoelectricity arises when certain crystalline materials develop an electric potential upon the application of mechanical stress, a behavior rooted in their asymmetric internal lattice structure. Quartz is a prominent example of such a material, as its unique silicon dioxide arrangement allows it to effectively interconvert electrical and mechanical energy. Consequently, it is the most widely utilized substance in technological applications requiring piezoelectric properties, such as in watches and various electronic sensors.

The reducing strength of phosphorus-based oxoacids is determined by the number of hydrogen atoms bonded directly to the central phosphorus atom. Within the structure of $H_3PO_2$, the presence of two P-H bonds is the critical factor that allows the molecule to act as a potent reducing agent.

The bond length between carbon atoms is fundamentally determined by the hybridization state, where lower s-character corresponds to longer bond distances. In diamond, carbon atoms undergo $sp^3$ hybridization and are connected exclusively by single $\sigma$-bonds, which are inherently longer than the bonds found in other carbon allotropes. In contrast, structures like graphite, $C_{60}$, and $C_{70}$ feature $sp^2$ hybridized carbon, which introduces partial double-bond character that significantly reduces the bond length.

The small atomic size of boron, combined with the absence of available d-orbitals, prevents $\mathrm{BCl}_3$ from dimerizing, whereas heavier group 13 trihalides such as $\mathrm{AlCl}_3$, $\mathrm{GaCl}_3$, and $\mathrm{InCl}_3$ readily form dimeric species like $\mathrm{Al}_2\mathrm{Cl}_6$ through chlorine bridging to complete their octets. As a result, $\mathrm{BCl}_3$ persists as a monomeric, electron-deficient Lewis acid that forms the adduct $\mathrm{Cl}_3\mathrm{B} \leftarrow \mathrm{NH}_3$ but lacks the structural capacity to participate in the self-dimerization observed in its heavier group counterparts.

Potassium is an alkali metal that consistently exhibits an oxidation state of +1 in its ionic compounds, which facilitates the calculation of the oxidation state for the oxygen atoms through algebraic charge balancing. In the compound $\mathrm{KO}_2$, solving the equation $1 + 2x = 0$ yields $x = -\frac{1}{2}$ for each oxygen atom, identifying the species as a superoxide.

The structure of $P_4O_6$ comprises a central $P_4$ tetrahedron with an oxygen atom inserted into each of the six $P-P$ edges. Since every phosphorus atom is bonded to three oxygen atoms, the total number of $P-O$ bonds in the molecule is calculated as $4 \times 3 = 12$.

The partial hydrolysis of xenon hexafluoride ($XeF_6$) occurs in sequential steps to produce xenon oxytetrafluoride ($XeOF_4$) and xenon dioxydifluoride ($XeO_2F_2$), described by the equations $XeF_6 + H_2O \rightarrow XeOF_4 + 2HF$ and $XeOF_4 + H_2O \rightarrow XeO_2F_2 + 2HF$. In both resulting compounds, the oxidation state of xenon remains $+6$ because oxygen carries a charge of $-2$ and fluorine carries $-1$, which balances the overall molecular charge without changing the oxidation state of the xenon atom from its initial $+6$ state \in $XeF_6$.