P Block Elements – NEET UG ChemistryPractice Questions & PYQs

Molecular geometry and Lewis structures dictate that single bonds consist entirely of sigma bonds, whereas multiple bonds incorporate additional pi bonds formed by the lateral overlap of unhybridized orbitals. Water ($\mathrm{H_2O}$) features an oxygen atom bonded to two hydrogen atoms exclusively through single covalent bonds, leaving it with two lone pairs and zero multiple bonds. In contrast, sulphur dioxide, nitrogen dioxide, and carbon dioxide all contain double bonds that inherently incorporate pi bonding components.

Thermal decomposition behavior of metal salts depends heavily on the redox characteristics of both their cationic and anionic components during heating. Potassium chlorate and zinc chlorate decompose to yield metal chlorides alongside gaseous oxygen, while potassium dichromate yields potassium chromate, chromium(III) oxide, and oxygen gas. Conversely, ammonium dichromate contains a reducing ammonium cation and an oxidizing dichromate anion that undergo internal redox reactions, producing nitrogen gas, chromium(III) oxide, and water vapor instead of releasing oxygen.

Interstitial compounds are formed when small non-metal atoms like hydrogen, carbon, or nitrogen become trapped inside the crystal lattices of transition metals. Because these inserted atoms do not alter the core metallic bonding network, these substances retain high melting points, extreme hardness, and metallic conductivity while remaining chemically inert. Consequently, describing them as chemically reactive directly contradicts their stable, unreactive nature.

Acidic strength among hydrogen compounds of group elements increases down the group due to a progressive decrease in bond dissociation enthalpy. Because $pK_a$ is inversely proportional to acid strength, stronger acids possess lower $pK_a$ values rather than higher ones. As hydrides progress from water to tellurium hydride, their acidity increases significantly, meaning their $pK_a$ values strictly decrease in that order. Thus, the sequence indicating an increase in $pK_a$ values for these group 16 hydrides is incorrect.

The basicity of oxoacids of phosphorus depends directly on the number of ionizable hydrogen atoms attached directly to oxygen atoms as $\mathrm{P}-\mathrm{OH}$ groups. Hypophosphorous acid ($\mathrm{H}_{3}\mathrm{PO}_{2}$) contains only one $\mathrm{P}-\mathrm{OH}$ bond alongside two non-ionizable $\mathrm{P}-\mathrm{H}$ bonds, making it a monoprotic acid rather than a diprotic one. All other provided statements accurately reflect the structural characteristics and industrial applications of phosphorus oxoacids.

Thermal decomposition of heavy metal sulfates typically yields a metal oxide and sulfur trioxide gas upon heating. When ferric sulfate ($\mathrm{Fe}_{2}(\mathrm{SO}_{4})_{3}$) is subjected to high temperatures, it thermally decomposes into solid ferric oxide ($\mathrm{Fe}_{2}\mathrm{O}_{3}$) and three moles of sulfur trioxide gas ($\mathrm{SO}_{3}$). The alternative choices involve reactions that produce sulfur dioxide or other reduction byproducts instead of pure sulfur trioxide.

Bond angles in central chlorine species are dictated by the number of lone pairs on the central atom and the steric repulsion between bonded substituents. Chlorine dioxide possesses a single unpaired electron and two lone pairs, resulting in a widened bond angle of approximately $118^\circ$, while dichlorine monoxide features two bulky chlorine atoms pressing against two lone pairs, compressing the bond angle down to roughly $111^\circ$. In the chlorite ion ($ClO_2^-$), the presence of two full lone pairs exerts maximum repulsive force on the bonding pairs, further squeezing the angle to around $103^\circ$ and establishing the correct increasing sequence.

The oxidation state in a neutral compound is determined by assigning standard oxidation numbers of $+1$ for hydrogen and $-2$ for oxygen, then solving the algebraic sum for phosphorus. For pyrophosphorous acid ($H_4P_2O_5$), setting up the equation $4(+1) + 2x + 5(-2) = 0$ yields $2x = 6$, giving $x = +3$. For hypophosphoric acid ($H_4P_2O_6$), the calculation $4(+1) + 2x + 6(-2) = 0$ results in $2x = 8$, leading to $x = +4$. Finally, for pyrophosphoric acid ($H_4P_2O_7$), evaluating $4(+1) + 2x + 7(-2) = 0$ gives $2x = 10$, which solves to $x = +5$.

The molecular architecture of phosphorus pentoxide is built upon a tetrahedral $P_4$ framework where each of the six edges connecting the phosphorus atoms contains a shared oxygen atom acting as a bridge. Additionally, each phosphorus atom holds one terminal double-bonded oxygen atom pointing outward from the cage structure. Counting these structural linkages reveals that exactly six oxygen atoms serve to bridge the phosphorus centers together within $P_4O_{10}$.

Oxidising strength across the halogen group depends heavily on thermodynamic factors including low bond dissociation energy and high hydration enthalpy of the resulting halide ions. Because fluorine possesses the highest positive standard reduction potential among the halogens, it has an immense thermodynamic tendency to accept electrons and undergo reduction. This exceptionally high reactivity makes $F_2$ capable of oxidizing all other halide species in solution, establishing it as the most powerful oxidising agent in the group.

Central oxygen in the ozone molecule undergoes $sp^2$ hybridization, forming two single covalent bonds with the terminal oxygen atoms. Each of these directional overlaps constitutes a standard $\sigma$ bond, while the remaining unhybridized p-orbitals overlap sideways to create a delocalized $\pi$ bond shared across the three centers. Consequently, the overall resonance hybrid structure accounts for a total of two $\sigma$ bonds and one $\pi$ bond.

Molecules adopt symmetrical geometries to equalize bond lengths unless lone pairs occupy specific positions that cause structural distortion. In $ClF_3$, the central chlorine atom accommodates three bond pairs and two lone pairs in a T-shaped geometry based on a trigonal bipyramidal arrangement. Because lone pair-lone pair and lone pair-bond pair repulsions compress the axial bonds differently than the equatorial ones, the axial bonds become significantly longer than the equatorial bonds.

Bond dissociation enthalpy normally decreases down a group as atomic size increases and overlap becomes less effective, but fluorine experiences an anomalous drop because of its exceptionally compact size. The extreme proximity of non-bonding electrons on the tiny fluorine atoms creates intense inter-electronic repulsion within the 2p orbitals, making the $F-F$ bond remarkably weak. Consequently, the actual bond dissociation energy sequence places chlorine ahead of fluorine, rendering the stated order incorrect.

Molecular dipole moment is determined by the vector sum of individual bond dipoles and the orbital dipole of the lone pair. In $NH_3$, nitrogen is less electronegative than hydrogen, causing the N-H bond dipoles to point towards the nitrogen atom in the exact same direction as the upward pull of the lone pair's orbital dipole. Conversely, fluorine is far more electronegative than nitrogen, which forces the N-F bond moments to point away from the nitrogen atom and directly oppose the direction of the lone pair dipole. This antagonistic alignment in $NF_3$ leads to substantial cancellation, resulting in a much smaller net dipole moment compared to $NH_3$.

The basic strength of hydrolyzable anions derived from group 16 elements diminishes as the size of the central atom increases down the group. Since basic character directly dictates the hydroxide ion concentration in isomolar salt solutions, higher basicity translates directly to a higher pH value. Therefore, oxide ion undergoes the most extensive hydrolysis to yield the highest alkalinity, while telluride ion produces the least alkaline solution.

A permanent dipole moment arises in molecules possessing an asymmetrical geometry where individual bond dipoles fail to cancel each other out. Among the choices, $\mathrm{SiF}_{4}$ and $\mathrm{BF}_{3}$ adopt symmetrical tetrahedral and trigonal planar shapes respectively, resulting in a net dipole moment of zero. Similarly, $\mathrm{XeF}_{4}$ features a square planar geometry with symmetrically opposed lone pairs that cancel out polarities. In contrast, $\mathrm{SF}_{4}$ undergoes $sp^{3}d$ hybridization resulting in a see-saw shape with a lone pair in the equatorial position, causing an asymmetric charge distribution and a net permanent dipole moment.

Paramagnetism arises exclusively in chemical species that contain one or more unpaired electrons in their molecular structures. While carbon dioxide, silicon dioxide, and sulfur dioxide possess completely paired electron configurations and exhibit diamagnetism, chlorine dioxide contains an odd number of total valence electrons. This odd electron configuration leaves a single unpaired electron delocalized over the molecule, causing it to be attracted to external magnetic fields.

The acidic strength of a series of oxoacids is fundamentally governed by the oxidation state of the central halogen atom. As the number of oxygen atoms bonded to the chlorine increases, the oxidation state of the central atom rises progressively from $+1$ in $\mathrm{HClO}$ up to $+7$ in $\mathrm{HClO}_4$. This escalation in positive charge enhances the electron-withdrawing inductive effect, which weakens the $\mathrm{O-H}$ bond and stabilizes the resulting conjugate base. Consequently, the tendency to donate a proton grows significantly, yielding the acidity sequence $\mathrm{HClO} < \mathrm{HClO}_2 < \mathrm{HClO}_3 < \mathrm{HClO}_4$.

Halogen bond dissociation enthalpies generally follow a decreasing trend down the group, yet fluorine displays an anomalously low bond energy owing to intense electrostatic repulsions between lone pairs on the small atoms. Silicon frequently expands its octet to attain a coordination number of six rather than four due to vacant $d$-orbitals. Furthermore, manganese maintains higher stability in the $\mathrm{Mn}(\mathrm{II})$ state compared to $\mathrm{Mn}(\mathrm{III})$ in aqueous environments because of its half-filled $d^5$ configuration, while group fifteen elements commonly exhibit $-3$, $+3$, and $+5$ oxidation states.

Phosphorus pentoxide acts as a powerful dehydrating agent and serves chemically as the anhydride of phosphoric acid. Upon reacting with water under heating conditions, it readily hydrates to form orthophosphoric acid as the sole product. The balanced reaction involves one mole of this oxide combining with three moles of water to yield two moles of $\mathrm{H}_{3} \mathrm{PO}_{4}$.