P Block Elements – JEE Main ChemistryPractice Questions & PYQs

Lithium aluminium hydride ($LiAlH_4$) functions as a potent source of hydride ions ($H^-$), which facilitate the reduction of silicon tetrachloride by substituting the chlorine atoms with hydrogen atoms to form silane ($SiH_4$). The residual lithium and aluminium species subsequently combine with the displaced chloride ions to yield lithium chloride ($LiCl$) and aluminium chloride ($AlCl_3$) as byproduct salts. The balanced chemical equation for this reduction process is:

$LiAlH_4 + SiCl_4 \rightarrow LiCl + AlCl_3 + SiH_4$

In graphite, each carbon atom undergoes $sp^2$ hybridisation involving one s-orbital and two p-orbitals, which results in a p-character percentage of $\frac{2}{3} \times 100 \approx 67\%$. Conversely, in diamond, each carbon atom exhibits $sp^3$ hybridisation formed by one s-orbital and three p-orbitals, leading to a p-character percentage of $\frac{3}{4} \times 100 = 75\%$.

This problem revolves around the amphoteric nature of aluminum, a property that allows it to react with both acids and strong bases. When aluminum metal reacts with aqueous $\mathrm{NaOH}$, it forms a white gelatinous precipitate of aluminum hydroxide, $\mathrm{Al(OH)_3}$. This precipitate exhibits amphoteric behavior, meaning it dissolves in the presence of excess $\mathrm{NaOH}$ to yield soluble sodium aluminate, $\mathrm{NaAlO_2}$. Upon strong heating, $\mathrm{Al(OH)_3}$ decomposes to produce $\mathrm{Al_2O_3}$, a stable oxide extensively employed as an adsorbent in chromatography techniques.

Many nitrogen-rich compounds, such as metal azides and specific ammonium salts, undergo thermal decomposition to liberate nitrogen gas because their molecular structures facilitate the formation of the stable $\mathrm{N_2}$ molecule. Barium azide, represented as $\mathrm{Ba(N_3)_2}$, decomposes to yield nitrogen along with metallic barium, while both $\mathrm{(NH_4)_2Cr_2O_7}$ and $\mathrm{NH_4NO_2}$ produce nitrogen gas through their respective decomposition reactions. Since $\mathrm{(NH_4)_2SO_2}$ lacks the chemical characteristics required to release nitrogen gas upon heating, it does not produce $\mathrm{N_2}$ like the other substances listed.

The structural bonding in nitrogen oxides determines the presence of direct nitrogen-nitrogen connectivity. Both $N_2O_3$ ($O=N-NO_2$) and $N_2O_4$ ($O_2N-NO_2$) contain a direct N-N covalent bond between the two nitrogen atoms. In contrast, $N_2O_5$ ($O_2N-O-NO_2$) features an oxygen bridge between the nitrogen centers, meaning it lacks a nitrogen-nitrogen bond.

The structure of $\mathrm{XeO_3F_2}$ is determined by the central Xenon atom forming bonds with three oxygen atoms and two fluorine atoms, utilizing all eight of its valence electrons. Xenon forms three double bonds with oxygen atoms and two single bonds with fluorine atoms, where each double bond consists of one $\sigma$-bond and one $\pi$-bond, leading to 3 $\pi$-bonds in total. Summing the $\sigma$-bonds from both oxygen and fluorine atoms results in $3 + 2 = 5$ bond pairs, and because all eight valence electrons are fully utilized in these bonds, there are 0 lone pairs on the central Xenon atom.

The identification of the chloride ion relies on its reaction with silver nitrate to form a characteristic white precipitate of silver chloride, $AgCl$, which is insoluble in dilute nitric acid. Sodium chloride, being a salt derived from a strong acid and a strong base, creates a neutral aqueous solution that does not alter the color of litmus paper. The reaction is represented by the equation $Ag^+ + Cl^- \rightarrow AgCl(s)$, where the resulting precipitate remains solid upon the addition of the acid, distinguishing it from other anions that would either dissolve with effervescence or produce a differently colored precipitate.

The molecular geometry and connectivity of sulfur oxoacids dictate the nature of their functional groups, which directly determines the count of double and single bonds around the sulfur centers. Both peroxodisulphuric acid ($H_2S_2O_8$) and pyrosulphuric acid ($H_2S_2O_7$) contain two sulfur atoms that maintain an identical local bonding environment. In both structures, every sulfur atom is double-bonded to two oxygen atoms and single-bonded to one hydroxyl group, meaning that even though the bridge connecting the two sulfur centers differs-being a peroxide linkage in the former and an oxygen bridge in the latter-the total count for each molecule remains four S=O bonds and two S-OH bonds.

Xenon hexafluoride ($XeF_6$) undergoes partial hydrolysis in the presence of water to form xenon oxytetrafluoride ($XeOF_4$) and hydrogen fluoride through the reaction $XeF_6 + H_2O \rightarrow XeOF_4 + 2HF$. When $XeF_6$ comes into contact with silica ($SiO_2$), it acts as a fluorinating agent to yield $XeOF_4$ and silicon tetrafluoride via the reaction $2XeF_6 + SiO_2 \rightarrow 2XeOF_4 + SiF_4$. As both reactions consistently produce the same product, compound 'X' is $XeOF_4$.

The structure of pyrophosphoric acid, $H_4P_2O_7$, contains a central $P-O-P$ linkage where each phosphorus atom is bonded to one terminal oxygen atom via a double bond and two hydroxyl ($-OH$) groups, resulting in a total of four $P-OH$ bonds per molecule. The oxidation state of the phosphorus atom is determined by assigning oxidation numbers of $+1$ for hydrogen and $-2$ for oxygen to the molecular formula, yielding the following calculation:
$4(+1) + 2(x) + 7(-2) = 0$
$2x - 10 = 0$
$x = +5$

The number of $\pi$-bonds in oxoacids of sulfur is determined by counting the $S=O$ double bonds present in their respective chemical structures. $H_2SO_3$ contains one $S=O$ bond, $H_2SO_4$ contains two $S=O$ bonds, and $H_2S_2O_7$ possesses four $S=O$ bonds, as it features two $S=O$ bonds on each sulfur atom connected by an oxygen bridge. Consequently, the decreasing order of $\pi$-bond count is $H_2S_2O_7 > H_2SO_4 > H_2SO_3$.

The reaction between chlorine gas and cold, dilute aqueous sodium hydroxide is a classic disproportionation process where chlorine acts as both an oxidizing and a reducing agent. As the chlorine atoms undergo simultaneous oxidation and reduction, their oxidation state shifts from zero to -1 and +1, resulting in the formation of chloride and hypochlorite ions. This chemical change is defined by the equation $\mathrm{Cl}_2 + 2\mathrm{OH}^- \rightarrow \mathrm{Cl}^- + \mathrm{ClO}^- + \mathrm{H}_2\mathrm{O}$.

Lithium is the only alkali metal that reacts directly with nitrogen at room temperature to form the nitride $Li_3N$, which follows the stoichiometry $M_3N$. This compound is thermally unstable at high temperatures, decomposing back into lithium metal and nitrogen gas, and it reacts with water to release ammonia, which is gas $B$. Ammonia gas subsequently reacts with aqueous $CuSO_4$ to form the soluble deep blue complex $[Cu(NH_3)_4]SO_4$, thereby identifying the metal $M$ as lithium and gas $B$ as $NH_3$.

Silica gel acts as a powerful adsorbent, making it the standard choice as a drying agent to protect sensitive goods from moisture. In the realm of electronics, silicon serves as the essential semiconductor material required to fabricate transistors and integrated circuits. Silicones are synthetic polymers renowned for their heat stability and water repellency, which makes them perfect for sealing gaps as sealants. Finally, complex structures known as silicates are frequently employed in chemical processes, particularly as ion-exchangers to remove hardness from water.

The electronic configuration $1s^{2}2s^{2}2p^{6}3s^{2}3p^6$ corresponds to the stable structure of Argon, which involves 18 electrons occupying orbitals up to the $3p$ subshell. Upon dissociation in an aqueous solution, sodium fluoride ($NaF$) yields $Na^+$ and $F^-$ ions, both of which possess the $1s^{2}2s^{2}2p^6$ configuration characteristic of Neon. Because these ions are isoelectronic with Neon and lack the $3s$ and $3p$ subshells entirely, the aqueous solution of this salt does not contain any ions with the specified Argon-like arrangement.

The gaseous $S_2$ molecule, much like dioxygen ($O_2$), possesses unpaired electrons in its antibonding orbitals, which gives rise to paramagnetic properties. Solid sulfur allotropes such as rhombic and monoclinic sulfur are composed of $S_8$ units arranged in a characteristic crown-like ring structure. Because the ring strain and geometric constraints differ between the eight-membered $S_8$ ring and the six-membered chair-shaped $S_6$ ring, the S-S-S bond angles are inherently different, meaning the claim that these angles are identical is factually incorrect.

Electrical conductivity in solids fundamentally depends on the availability of mobile charge carriers rather than merely the hybridization state of the constituent atoms. Diamond consists of a rigid, three-dimensional tetrahedral lattice where every carbon atom is $sp^3$ hybridized, locking all valence electrons into localized sigma bonds and preventing current flow. Graphite possesses a planar structure where carbon atoms are $sp^2$ hybridized, allowing the remaining unhybridized $p$-orbitals to form a delocalized electron cloud that facilitates conductivity; therefore, while the hybridization descriptions are factually accurate, they describe the bonding geometry and do not serve as the complete mechanistic explanation for why one allotrope conducts electricity and the other does not.

The oxidation state of a phosphorus atom within its oxoacids is fundamentally determined by the arrangement of atoms, with the +3 state defining the characteristic chemistry of the phosphorous acid group. Both orthophosphorous acid, expressed as $\mathrm{H_3PO_3}$, and pyrophosphorous acid, defined by the formula $\mathrm{H_4P_2O_5}$, contain phosphorus atoms specifically in this +3 oxidation state. Recognizing these two particular acids identifies the correct pair that satisfies the requirement for a +3 formal oxidation state.

Electronegativity defines an atom's capacity to pull shared electrons toward its nucleus, meaning the element with the highest electronegativity always gains electron density rather than losing it in a bond. Because fluorine sits at the peak of the electronegativity scale in the periodic table, it dominates every covalent bond by attracting shared electrons, preventing any other atom from withdrawing them. This behavior forces the element to be restricted to an oxidation state of -1 across all of its chemical combinations, which makes it impossible for fluorine to ever assume a positive value.

Xenon consists of individual, nonpolar atoms with a stable valence shell, meaning it does not naturally possess a permanent dipole or charge. As a result, the only way these atoms can experience attractive forces is through the temporary, random fluctuations in their electron distribution that create instantaneous dipole-induced dipole forces. These weak, fleeting attractions are the primary mechanism enabling xenon gas to overcome thermal energy and condense into the liquid phase.