Some P Block Elements – NEET UG ChemistryPractice Questions & PYQs

The acidic strength of hydrogen halides is determined by the ease with which the H-X bond undergoes dissociation, which is directly governed by the bond enthalpy. As the atomic size of the halogen increases down the group, the orbital overlap between hydrogen and the halogen becomes increasingly inefficient, resulting in a progressively weaker H-X bond. Consequently, the strength of the acid increases from $HF$ to $HI$, meaning the actual order is $HF < HCl < HBr < HI$, which renders the decreasing order provided in the option inherently incorrect.

Dilute sulphuric acid acts as a diagnostic reagent in salt analysis by displacing weaker acids from their salts, which results in the release of characteristic gases that reveal the identity of the parent anions. This process relies on distinct physical observations-such as gas colour or odour-that occur immediately upon treatment with the acid.

Carbonate salts ($\mathrm{CO_3^{2-}}$) react to liberate carbon dioxide, which is perceived as a brisk effervescence of a colourless gas. Sulphide salts ($\mathrm{S^{2-}}$) evolve hydrogen sulphide, easily recognized by its sharp smell of rotten eggs. Treating sulphite salts ($\mathrm{SO_3^{2-}}$) with the acid releases sulphur dioxide, which is identified by its pungent odour. Finally, nitrite salts ($\mathrm{NO_2^-}$) produce nitrogen dioxide upon decomposition, appearing as characteristic reddish-brown fumes. These reactions establish that carbonate corresponds to II, sulphide to III, sulphite to IV, and nitrite to I.

Systematic qualitative analysis organizes metal ions into distinct analytical groups based on their selective precipitation reactions with specific reagents. By identifying the analytical group for each cation-$\mathrm{Cu}^{2+}$ in Group II, $\mathrm{Al}^{3+}$ in Group III, $\mathrm{Co}^{2+}$ in Group IV, $\mathrm{Ba}^{2+}$ in Group V, and $\mathrm{Mg}^{2+}$ in Group VI-we can arrange them in ascending order of their group numbers. Aligning these ions with their corresponding identifiers leads to the sequence of $\mathrm{Cu}^{2+}$ (B), $\mathrm{Al}^{3+}$ (A), $\mathrm{Co}^{2+}$ (D), $\mathrm{Ba}^{2+}$ (C), and $\mathrm{Mg}^{2+}$ (E), resulting in the final order of B, A, D, C, E.

Group 16 elements typically feature an outer electronic configuration of $ns^2 np^4$, providing a strong inclination to gain two electrons to achieve a stable octet and manifest a -2 oxidation state. However, metallic character increases significantly moving down the group, which diminishes the ability of these elements to act as electron acceptors and form negative ions. Consequently, while Oxygen, Selenium, and Tellurium are known to exhibit the -2 oxidation state, the highly metallic Polonium is restricted to displaying only +2 and +4 oxidation states.

The allotropes of carbon exhibit unique structural and chemical characteristics that dictate their industrial utility. Coke serves as a vital reducing agent in metallurgical processes, while diamond is defined by its robust three-dimensional network where every carbon atom undergoes $\text{sp}^3$ hybridization, resulting in a rigid tetrahedral arrangement. Fullerene represents a distinct structural class as a cage-like molecule composed of a specific configuration of five-membered and six-membered rings. Finally, graphite possesses a layered structure that makes it naturally soft and slippery, allowing it to function effectively as a dry lubricant in heavy machinery operating at high temperatures.

Hydrides are primarily categorized based on the nature of their bonding and the distribution of their valence electrons relative to the octet rule. Sodium hydride ($\mathrm{NaH}$) is a classic ionic or saline hydride, whereas phosphine ($\mathrm{PH}_3$) acts as an electron-rich hydride due to the presence of lone pairs on the central atom. Germane ($\mathrm{GeH}_4$) is classified as electron-precise since it utilizes four valence electrons to form four bonds, and $\mathrm{LaH}_{2.87}$ serves as a typical example of a non-stoichiometric or metallic hydride.

Lewis acids are substances capable of accepting a lone pair of electrons because they possess vacant orbitals or an incomplete octet in their valence shell. Within this group, $\mathrm{BF}_3$ functions as a Lewis acid because the central boron atom is electron-deficient and able to accommodate additional electron density. In contrast, species such as $\mathrm{H}_2\mathrm{O}$, $\mathrm{NH}_3$, and $\mathrm{OH}^{-}$ contain available lone pairs that they can donate, which characterizes them as Lewis bases.

As you descend Group 13, the poor shielding effect of intervening $d$ and $f$ orbitals increases the effective nuclear charge, which binds the outer $ns$ electrons tightly and inhibits their participation in chemical bonding. This phenomenon, known as the inert pair effect, ensures that the +1 oxidation state becomes increasingly stable relative to the +3 state for heavier elements. This trend is demonstrated by the reduction potentials, where the value of $E^{\circ} (\mathrm{Tl}^{3+} \mid \mathrm{Tl}^{+}) = +1.6\text{ V}$ is significantly higher than the value of $E^{\circ} (\mathrm{In}^{3+} \mid \mathrm{In}^{+}) = -0.16\text{ V}$. Because $\mathrm{Tl}^{3+}$ acts as a powerful oxidizing agent that effectively reduces to $\mathrm{Tl}^{+}$, the compound $\mathrm{TlI}$ is considerably more stable than $\mathrm{TlI}_3$.

The structure of diborane is characterized by electron-deficient bonding, where each boron atom adopts $sp^3$ hybridization to facilitate the formation of both terminal and bridging bonds. These orbitals allow the boron atoms to engage in the distinctive 3-centre-2-electron bridge bonds while maintaining four standard 2-centre-2-electron terminal B-H bonds within a planar framework. Consequently, claiming that the boron atoms are $sp^2$ hybridized is inaccurate because their actual bonding environment requires a tetrahedral geometry, which contradicts the planar characteristics typically associated with $sp^2$ hybridization.

Thermal decomposition is the chemical process where a compound breaks down into simpler substances upon the application of high heat. In the case of Borax, this process yields two primary products: sodium metaborate and the glassy, transparent substance known as boric anhydride. The reaction is represented by $\mathrm{Na}_2 \mathrm{B}_4 \mathrm{O}_7 \xrightarrow{\text{heat}} \mathrm{B}_2 \mathrm{O}_3 + 2 \mathrm{NaBO}_2$, which confirms that the product X is $\mathrm{B}_2 \mathrm{O}_3$.

The hybridization state of carbon in its allotropes depends fundamentally on the geometry of its covalent bonding. In diamond, each carbon atom establishes four sigma bonds with its neighbors to form a rigid, three-dimensional lattice, which forces the carbon atoms into an $sp^3$ hybridized state. In contrast, the structure of graphite consists of two-dimensional hexagonal layers where each carbon atom binds to only three neighbors, leaving one electron in a $p$-orbital for pi bonding, thereby resulting in an $sp^2$ hybridized state.

Understanding the chemical composition of these boron-based compounds relies on identifying their specific hydration levels and derivative forms. Borax is a decahydrate represented by the formula $\mathrm{Na}_2 \mathrm{B}_4 \mathrm{O}_7 \cdot 10 \mathrm{H}_2 \mathrm{O}$, whereas kernite is its tetrahydrate counterpart, $\mathrm{Na}_2 \mathrm{B}_4 \mathrm{O}_7 \cdot 4 \mathrm{H}_2 \mathrm{O}$. Orthoboric acid is the familiar $\mathrm{H}_3 \mathrm{BO}_3$, and the glass-like borax bead produced during chemical analysis is defined as sodium metaborate, $\mathrm{NaBO}_2$.

Synthesis gas is the term defined by the mixture of carbon monoxide and hydrogen gas, expressed as $CO(g)+H_2(g)$. Temporary hardness of water arises from the presence of $Ca(HCO_3)_2+Mg(HCO_3)_2$, whereas diborane, $B_2H_6$, serves as a classic example of an electron deficient hydride. Hydrogen peroxide, $H_2O_2$, completes the set as a non-planar molecule characterized by an open-book geometry. These associations yield the sequence where (a) links to (iii), (b) corresponds to (i), (c) is paired with (ii), and (d) matches (iv).

Refrigeration for frozen goods relies on solid carbon dioxide, $CO_2(s)$, widely known as dry ice, which contradicts the claim that the gaseous form is utilized for this purpose. The molecular structure of $C_{60}$ is uniquely defined by twenty six-membered rings and twelve five-membered rings, rendering the second statement incorrect because it misidentifies these ring counts. In contrast, ZSM-5 functions as a critical zeolite catalyst for converting alcohols into gasoline, and carbon monoxide is recognized as a colorless and odorless gas, confirming the accuracy of both the final two statements.

Carbon monoxide is notoriously dangerous because it binds to hemoglobin in the bloodstream much more tightly than oxygen does. When this gas enters the body, it forms carboxyhemoglobin, a compound that is about 300 times more stable than oxyhemoglobin. This extremely high affinity severely inhibits the ability of the blood to transport oxygen, which is a lethal consequence frequently resulting from the incomplete combustion of fuels. Because carboxyhemoglobin remains bound to hemoglobin with such high stability, the claim that it is less stable than oxyhemoglobin is scientifically incorrect.

Although silicon possesses vacant d-orbitals that theoretically allow it to expand its coordination number beyond four, the physical size of the silicon atom imposes a critical geometric constraint. Specifically, $[SiCl_6]^{2-}$ cannot exist because the central silicon ion is too small to accommodate six bulky chloride ions around it without causing insurmountable steric repulsion. This is not an issue for species like $[SiF_6]^{2-}$, $[GeCl_6]^{2-}$, and $[Sn(OH)_6]^{2-}$, where the central atoms are appropriately sized or the ligands are small enough to allow for stable hexacoordination.

Silicones are synthetic organosilicon polymers characterized by the repeating unit $(R_2SiO)_n$ that exhibit exceptional biocompatibility, chemical inertness, and thermal stability. These properties make them the material of choice for cosmetic surgery, particularly in the production of breast implants and soft tissue fillers, as they are non-toxic and do not trigger significant adverse immune responses within the human body. Unlike silica, silicates, or zeolites, silicones possess the elasticity and structural integrity required for successful integration into medical surgical procedures.

Group 14 tetrahalides typically feature covalent bonding, but the nature of the interaction shifts significantly when the central metal bonds with highly electronegative halogens like fluorine. Both $SnF_4$ and $PbF_4$ are categorized as ionic compounds rather than covalent ones, as the extreme electronegativity of fluorine dictates strong ionic character within these structures. Describing $PbF_4$ as covalent in nature contradicts this established ionic property, making that statement the incorrect one.

The ability of an element to form complex ions like $MF_6^{3-}$ depends fundamentally on its maximum covalency, which is determined by the presence of vacant orbitals. Boron is situated in the second period of the periodic table, meaning its valence shell contains only 2s and 2p orbitals, leaving no vacant d-orbitals available for bonding. Consequently, boron is strictly limited to a maximum coordination number of 4 and cannot expand its octet to accommodate six fluorine atoms. In contrast, heavier group 13 elements like aluminum, gallium, and indium possess accessible d-orbitals that allow them to extend their covalency and successfully form the $MF_6^{3-}$ ion.

The inert pair effect describes the tendency of $ns^2$ valence electrons in heavier p-block elements to remain non-bonding, a phenomenon that stabilizes the +2 oxidation state for lead while keeping the +4 oxidation state more stable for tin. Because $Sn^{2+}$ readily loses electrons to reach the more stable $Sn^{4+}$ configuration, it functions as a reducing agent, whereas $Pb^{4+}$ tends to gain electrons to revert to the more stable $Pb^{2+}$ state, thereby acting as an oxidising agent.