P Block Elements – JEE Main ChemistryPractice Questions & PYQs

Phosphorus allotropes exhibit fundamentally different chemical behaviors when subjected to alkaline conditions, which dictates the structure of the resulting oxoacids. While white $P_4$ reactions typically generate products containing $P-H$ bonds, red $P_4$ reacts with alkali to produce hypophosphoric acid, represented as $H_4 P_2 O_6$. This particular oxoacid is characterized by the absence of $P-H$ bonds, making red $P_4$ the specific reactant $A$ required by the reaction conditions.

White phosphorus undergoes a disproportionation reaction in the presence of strong alkalis because it acts simultaneously as an oxidizing and a reducing agent. When heated with concentrated sodium hydroxide solution, the phosphorus molecules are transformed into phosphine gas and sodium hypophosphite, as described by the reaction:

$\mathrm{P}_4 + 3\mathrm{NaOH} + 3\mathrm{H}_2\mathrm{O} \rightarrow 3\mathrm{NaH}_2\mathrm{PO}_2 + \mathrm{PH}_3$

The efficiency of industrial chemical processes often relies on selecting specific catalysts to lower activation energy and achieve high selectivity for desired products. For the synthesis of ammonia via the reaction $\mathrm{N}_2(g) + 3\mathrm{H}_2(g) \rightarrow 2\mathrm{NH}_3(g)$, a multi-component heterogeneous catalyst composed of $\mathrm{Fe}_x\mathrm{O}_y + \mathrm{K}_2\mathrm{O} + \mathrm{Al}_2\mathrm{O}_3$ is essential to drive the Haber process forward. When hydrogenating carbon monoxide, the choice of catalyst determines the specific product formed: the methanation reaction $\mathrm{CO}(g) + 3\mathrm{H}_2(g) \rightarrow \mathrm{CH}_4(g) + \mathrm{H}_2\mathrm{O}(g)$ is promoted by the use of $\mathrm{Ni}$, while the production of formaldehyde $\mathrm{CO}(g) + \mathrm{H}_2(g) \rightarrow \mathrm{HCHO}(g)$ requires metallic $\mathrm{Cu}$ as the catalyst. Finally, the synthesis of methanol through the reaction $\mathrm{CO}(g) + 2\mathrm{H}_2(g) \rightarrow \mathrm{CH}_3\mathrm{OH}(g)$ is made possible by employing a specialized catalyst system consisting of $\mathrm{Cu}/\mathrm{ZnO}-\mathrm{Cr}_2\mathrm{O}_3$. These distinct catalytic materials align with the corresponding processes, ensuring that reactants successfully transform into their respective target molecules.

Diborane ($\mathrm{B}_2\mathrm{H}_6$) features a non-planar geometry with two bridging 3-centre-2-electron bonds and four terminal 2-centre-2-electron bonds, which results in non-equivalent B-H bonds. Because the boron atoms possess an incomplete octet, the compound is electron-deficient and behaves as a Lewis acid. It can be prepared via laboratory synthesis using common reagents like $\mathrm{BF}_3$ and $\mathrm{NaBH}_4$, such as in the reactions $2\,\mathrm{BF}_3+6\,\mathrm{NaH} \xrightarrow{450\,\mathrm{K}} \mathrm{B}_2\mathrm{H}_6+6\,\mathrm{NaF}$ and $2\,\mathrm{NaBH}_4+\mathrm{I}_2 \rightarrow \mathrm{B}_2\mathrm{H}_6+2\,\mathrm{NaI}+\mathrm{H}_2$

Chemical substances exhibit distinct biological and physiological impacts based on their unique chemical properties. While high levels of fluoride are clinically linked to the bending of bones, sulphates are well-documented for their laxative effect on the digestive system. Nicotine serves as a powerful pesticide to manage insect populations, whereas sodium arsinite is utilized specifically as a herbicide to control unwanted plant growth. Mapping these functions to their respective substances results in the pairing A-III, B-II, C-I, and D-IV.

When silver chloride encounters an aqueous ammonia solution, it dissolves because ammonia molecules act as ligands that coordinate with silver ions to create a highly stable, soluble complex. During this chemical transformation, each silver ion from the precipitate binds with two molecules of ammonia to form the $[\mathrm{Ag}(\mathrm{NH}_3)_2]^+$ cation, which effectively results in the soluble complex $[\mathrm{Ag}(\mathrm{NH}_3)_2] \mathrm{Cl}$. This entire reaction is defined by the chemical equation $\mathrm{AgCl} + 2 \mathrm{NH}_3 \rightarrow [\mathrm{Ag}(\mathrm{NH}_3)_2]^+ \mathrm{Cl}^-$.

The diagonal relationship between beryllium and aluminum leads to shared chemical characteristics, such as the formation of chlorine-bridged structures in their respective chlorides. While these chlorides are soluble in organic solvents, they act as Lewis acids due to the electron-deficient state or presence of vacant orbitals in the metal atoms, which makes the claim that they function as Lewis bases incorrect.

Furthermore, beryllium and aluminum exhibit amphoteric properties, allowing them to react with both acidic and alkaline environments. Consequently, their hydroxides dissolve readily when treated with an excess of alkali to produce the corresponding beryllate $[\mathrm{Be}(\mathrm{OH})_4]^{2-}$ and aluminate $[\mathrm{Al}(\mathrm{OH})_4]^{-}$ ions, confirming that the second statement is true.

The formation of oxoacids depends heavily on the atomic radius and electronegativity of the halogen involved. Since fluorine has the smallest atomic radius and the highest electronegativity among all halogens, its bonding capabilities are significantly limited, restricting it to the formation of only one stable oxoacid, $HOF$. Therefore, these unique physical characteristics of fluorine directly account for why it cannot form the range of oxoacids observed in other halogens, making the reason a valid explanation for the assertion.

The reaction of bromine with an excess of fluorine yields bromine pentafluoride, as represented by the balanced chemical equation $\mathrm{Br}_2 + 5\mathrm{F}_2 \rightarrow 2\mathrm{BrF}_5$. Within this specific interhalogen compound, the bromine atom exists in a +5 oxidation state. This oxidation state is the defining feature of the halate class of compounds, where the central halogen atom also maintains a +5 oxidation state, making it the category consistent with this reaction product.

The stability of nitrogen trihalides is primarily governed by the strength of the $N - X$ bond and the steric environment around the central nitrogen atom. Because nitrogen is a very small atom, it struggles to accommodate larger halogen atoms, leading to significantly decreased stability as you move down the group from fluorine to iodine. Consequently, $NF_3$ stands out as the most stable trihalide because the small fluorine atoms minimize steric crowding and allow for a stronger, more compact bond compared to the larger, bulkier halides.

When white phosphorus reacts with an aqueous alkali solution, it undergoes a disproportionation reaction to generate phosphine and a salt of hypophosphorous acid. The resulting oxoacid is $H_3PO_2$, which is systematically named phosphinic acid. This molecule contains a central phosphorus atom directly bonded to two hydrogen atoms, resulting in the two P-H bonds described.

The oxidation states of sulphur within oxoacids are determined by their specific structural geometry, where the connectivity of atoms often results in non-equivalent environments for the sulphur centres. In the case of $\mathrm{H}_2\mathrm{S}_2\mathrm{O}_3$, one sulphur atom occupies a central position bonded to oxygen atoms while the other acts as a terminal sulphur atom, leading to two distinct oxidation states of $+6$ and $-2$. Conversely, compounds such as $\mathrm{H}_2\mathrm{S}_2\mathrm{O}_6$, $\mathrm{H}_2\mathrm{S}_2\mathrm{O}_7$, and $\mathrm{H}_2\mathrm{S}_2\mathrm{O}_8$ maintain symmetrical structures in which both sulphur atoms are chemically equivalent, resulting in identical oxidation states for both atoms within those molecules.

Boric acid is classified as a weak acid because it does not function as a typical Arrhenius acid that directly releases protons on its own. Instead, it acts as a Lewis acid in an aqueous solution by accepting a hydroxide ion from water to form the $[\text{B}(\text{OH})_4]^-$ complex, which leads to the release of a proton according to the reaction $\text{H}_3\text{BO}_3 + \text{H}_2\text{O} \rightleftharpoons [\text{B}(\text{OH})_4]^- + \text{H}^+$. Because this indirect mechanism is the specific cause of its limited ability to donate protons, both the assertion regarding its classification as a weak acid and the explanation detailing its unique interaction with water are factually accurate and causally connected.

The structural complexity of organosilicon products is determined by the number of hydroxyl groups present on the silicon atom, which dictates how the molecules condense and polymerize. For instance, $\mathrm{(CH_3)_4Si}$ functions as a simple silane because it lacks hydroxyl groups to facilitate polymerization, while $\mathrm{(CH_3)_3Si(OH)}$ is limited to forming a dimeric silicone since it contains only one reactive hydroxyl group. Extending this principle, $\mathrm{(CH_3)_2Si(OH)_2}$ undergoes linear condensation to form chain silicones, whereas the presence of three hydroxyl groups in $\mathrm{(CH_3)Si(OH)_3}$ promotes multi-directional cross-linking, resulting in 2D-silicone structures.

The acidic strength of oxides is directly proportional to the oxidation state of the central atom, as a higher positive charge enhances the element's electronegative influence and increases its acidity. Therefore, the pentavalent oxide $E_2O_5$ is more acidic than the trivalent oxide $E_2O_3$, which makes the first statement false.

Furthermore, the acidic character of oxides within a group is determined by the non-metallic nature of the central element. As one proceeds down group 15, the non-metallic character of the elements diminishes, which directly leads to a decrease in the acidic character of the trivalent oxides $E_2O_3$, confirming that the second statement is true.

The reaction between $\mathrm{PCl}_3$ and water occurs via the hydrolysis of the phosphorus-chlorine bonds, where chlorine atoms are systematically replaced by hydroxyl groups. Through the sequential addition of water, the chlorine atoms are substituted to form hydroxy-derivatives as intermediate products, resulting in a species $\mathrm{B}$ that contains two hydroxyl groups attached to the phosphorus atom. Since each hydroxyl group provides one ionisable hydrogen atom and the bonds between phosphorus and the remaining chlorine or hydrogen do not dissociate in this context, the structure contains exactly 2 ionisable protons.

Atmospheric pollutants containing nitrogen play a critical role in environmental chemistry, where nitrogen dioxide ($\mathrm{NO}_2$) stands out due to its specific harmful impact on vegetation. When plants are exposed to high concentrations of $\mathrm{NO}_2$, it directly interferes with their metabolic functions, effectively hindering photosynthesis and causing significant physical degradation to their leaves.

Metalloids are elements defined by physical and chemical properties that sit between those of metals and non-metals. Among the choices provided, tellurium is a metalloid because it is situated along the zig-zag boundary line on the periodic table that distinguishes metallic and non-metallic character. Conversely, scandium functions as a transition metal, while lead and bismuth are classified as post-transition metals, meaning they exhibit distinct metallic behavior rather than the intermediate characteristics of a metalloid.

Tooth enamel consists primarily of the mineral fluoroapatite, which is chemically represented by the formula $3\mathrm{Ca}_3(\mathrm{PO}_4)_2 \cdot \mathrm{CaF}_2$. This structural composition incorporates calcium ions as $\mathrm{Ca}^{2+}$, fluoride as $\mathrm{F}^{-}$, and phosphorus exclusively in the $\mathrm{P}^{5+}$ oxidation state within the phosphate groups. Because the phosphorus atoms are restricted to these phosphate units within the crystal lattice, the $\mathrm{P}^{3+}$ ionic form is not present.

The atmospheric formation of acid rain is primarily driven by the accumulation of acidic oxides, especially those of sulfur and nitrogen, which react with environmental moisture. Sulfur dioxide, a key pollutant, undergoes a process involving atmospheric oxygen and water vapor to ultimately yield sulfuric acid. The complete chemical reaction describing this pathway to acid rain is $2\,\mathrm{SO}_2 + \mathrm{O}_2 + 2\,\mathrm{H}_2\mathrm{O} \rightarrow 2\,\mathrm{H}_2\mathrm{SO}_4$, which effectively captures how these reactants combine to create the acidic precipitation.