P Block Elements – JEE Main ChemistryPractice Questions & PYQs

Nitrogen molecules are chemically inert under ambient conditions due to the exceptionally strong triple bond ($N \equiv N$) that requires high bond dissociation energy to break. The reaction between nitrogen and oxygen to form nitric oxide, represented by $N_2(g) + O_2(g) \rightleftharpoons 2NO(g)$, is highly endothermic and possesses a significant activation energy barrier that can only be overcome at very high temperatures, such as those generated by lightning or internal combustion engines. Because standard atmospheric temperatures are insufficient to supply this activation energy, these gases do not react spontaneously, confirming that the specific temperature requirement is the causal factor for the stability of their mixture in the atmosphere.

Interhalogen compounds are generally more reactive than pure diatomic halogen molecules because of the specific nature of the bond shared between the atoms. In a compound such as $ICl$, the difference in electronegativity between the two different halogen atoms creates a polar covalent bond, whereas elemental halogens like $Cl_2$, $Br_2$, and $I_2$ contain non-polar covalent bonds. This polarity makes the bond within the interhalogen molecule easier to break during chemical reactions, rendering it more reactive than the corresponding pure halogen molecules.

The boiling points of noble gases are determined by the strength of London dispersion forces, which intensify as atomic size and polarizability increase down the group. Consequently, Xenon possesses the highest boiling point at 165 K among the elements listed, significantly exceeding the boiling points of Krypton at 119.7 K, Neon at 27.1 K, and Helium at 4.2 K. This trend demonstrates that larger atoms require greater thermal energy to overcome intermolecular attractions, thereby raising their boiling points.

Nitric oxide, NO, contains a total of 15 electrons, and according to molecular orbital theory, its electronic configuration leaves one unpaired electron in an antibonding $\pi^*$ orbital. The presence of this lone, unpaired electron makes the molecule paramagnetic in the gaseous state, rendering the statement that it is diamagnetic factually incorrect. Meanwhile, the molecule remains a neutral oxide with a bond order of 2.5 and readily reacts with oxygen to form nitrogen dioxide, confirming that the other listed properties are accurate.

As an alkali metal, cesium is restricted to a $+1$ oxidation state, rendering the existence of a $Cs^{3+}$ state impossible. The structure of $CsI_3$ is instead stabilized by the formation of the $Cs^+$ cation and the stable polyatomic $I_3^-$ anion. Models suggesting the presence of discrete $I_2$ molecules are invalid because those molecules are simply too large to be accommodated within the crystal lattice.

The acidity of oxoacids is fundamentally determined by the oxidation state of the central chlorine atom, where higher oxidation numbers promote greater electron withdrawal from the O-H bond. As the oxidation state of chlorine decreases from +7 in $HClO_4$ and +5 in $HClO_3$ to +3 in $HClO_2$ and +1 in $HOCl$, the tendency for the O-H bond to break and release a proton diminishes. Consequently, the strength of these acids follows the decreasing order $HClO_4 > HClO_3 > HClO_2 > HOCl$.

The physical and chemical properties of ozone, including its structure and state, are determined by the arrangement of its valence electrons and its molecular geometry. Ozone exists as a bent molecule, and in its condensed solid phase, it displays a characteristic violet-black color. Furthermore, as a gaseous molecule, ozone exhibits diamagnetic properties, which indicates that all of its electrons are paired. When evaluating whether species are isoelectronic, one must compare their total electron counts. In this case, calculations show that $ONCl$ contains $32$ electrons, whereas the $ONO^-$ ion contains $24$ electrons, which confirms that they possess different electron totals and are therefore not isoelectronic.

The bond angles in these pyramidal $\mathrm{AX}_3$ molecules decrease down the group due to a reduction in the repulsion between bond pairs as the central atom increases in size and bond lengths grow. Moving down from Nitrogen to Antimony, the atomic size expands and the electronegativity of the central atom decreases, which causes the bond angles to tighten progressively. This trend is demonstrated by values such as 107 degrees for $\mathrm{NCl}_3$, 94 degrees for $\mathrm{PCl}_3$, and 92 degrees for $\mathrm{AsCl}_3$, confirming that $\mathrm{SbCl}_3$ possesses the smallest bond angle in the series.

The inability of Boron to form $\mathrm{BF}_6^{3-}$ stems from its position in the second period of the periodic table, where its valence shell is restricted to $2s$ and $2p$ orbitals. Since Boron lacks available $d$-orbitals, it cannot expand its octet to accommodate a coordination number greater than four. Consequently, while stable species like $\mathrm{BH}_4^{-}$, $\mathrm{B(OH)}_4^{-}$, and $\mathrm{BO}_2^{-}$ satisfy the octet requirements, forming a hexacoordinate anion like $\mathrm{BF}_6^{3-}$ remains chemically impossible.

The stability of group 15 hydrides depends directly on the strength of the bond formed between the central atom and the hydrogen atom. As we descend the group, the atomic size of the central element increases, leading to a progressively poorer orbital overlap between that large central atom and the small hydrogen atom. This mismatch weakens the bond, causing stability to decrease rapidly from $\mathrm{NH}_3$ down to $\mathrm{BiH}_3$, which makes the claim that stability increases in this direction factually wrong.

The chemical behavior of sulphur is defined by its capacity to exist in oxidation states spanning the range from -2 to +6 within its diverse array of chemical compounds. Molecules such as hydrogen sulphide, $H_2S$, demonstrate that sulphur frequently adopts a -2 oxidation state, which is significantly lower than +4. Therefore, the assertion that the oxidation state of sulphur is never less than +4 is inaccurate, as it fails to account for the common lower oxidation states exhibited by the element.

The hydrolysis of xenon hexafluoride is a well-established chemical process represented by the equation $\mathrm{XeF}_6 + 3 \mathrm{H}_2\mathrm{O} \rightarrow \mathrm{XeO}_3 + 6 \mathrm{HF}$. In this interaction, the formation of $\mathrm{XeO}_3$ and $\mathrm{HF}$ is the thermodynamically favored outcome under standard conditions. Consequently, the reverse reaction, represented by $\mathrm{XeO}_3 + 6 \mathrm{HF} \rightarrow \mathrm{XeF}_6 + 3 \mathrm{H}_2\mathrm{O}$, is not feasible because the chemical system naturally proceeds toward the hydrolysis products rather than the reactants.

In the solid phase, both $\mathrm{BeCl}_2$ and $\mathrm{AlCl}_3$ stabilize their structures by forming extended polymeric networks where chloride atoms function as bridges between the central metal atoms. Boric acid is characterized as a Lewis acid because it acts as an electron pair acceptor when reacting with a hydroxide ion, which prevents it from being classified as a protonic acid. Furthermore, the small atomic radius of beryllium limits its capacity to accommodate ligands, restricting its coordination number to 4 rather than 6. Consequently, the statement regarding $\mathrm{B}_2 \mathrm{H}_6 \cdot 2 \mathrm{NH}_3$ is inaccurate, as the presence of bridged chloride structures represents the only valid chemical description among the choices.

Silicone polymerization is dictated by the functionality of the precursor monomer, where each chlorine atom provides a reactive site for the formation of siloxane linkages during hydrolysis. Among the choices, $\mathrm{RSiCl}_3$ possesses three chlorine atoms, which allows the resulting intermediate to participate in condensation reactions in three distinct directions simultaneously. This trifunctionality causes the growing chains to branch and interconnect extensively, creating a rigid, three-dimensional cross-linked network rather than the simple linear polymers typical of precursors with fewer reactive sites.

The reaction between chlorine and ammonia is determined by the stoichiometric ratio of the reactants. When chlorine interacts with an excess of ammonia, the hydrogen chloride formed during the reaction is immediately neutralized by the excess ammonia, resulting in the quantitative formation of ammonium chloride instead. Therefore, the final products are nitrogen gas and ammonium chloride ($\mathrm{NH}_4\mathrm{Cl}$), which makes the statement describing the products as $\mathrm{N}_2$ and $\mathrm{HCl}$ factually incorrect.

The stability of group 14 dihalides is determined by the inert pair effect, which describes the growing reluctance of valence s-electrons to engage in covalent bonding due to inefficient shielding by inner orbitals as atomic size increases down the group. Because these electrons become less available for bonding in heavier elements, the +2 oxidation state becomes progressively more stable relative to the +4 oxidation state as we descend the group. This trend results in a steady increase in stability across the series $\mathrm{SiX}_2 < \mathrm{GeX}_2 < \mathrm{SnX}_2 < \mathrm{PbX}_2$.

Ammonium sulphate fertilizers contribute to soil acidification due to the chemical behavior of the compound when it encounters moisture. Upon reacting with water in the soil, it follows the reaction $\mathrm{(NH_4)_2SO_4}+2\mathrm{H_2O}\rightarrow 2\mathrm{H_2SO_4}+\mathrm{NH_4OH}$. Because this reaction produces $\mathrm{H_2SO_4}$, which is a strong acid, regular application results in a progressive increase in the soil's acidity.

Disproportionation is a chemical phenomenon where a single species undergoes both oxidation and reduction simultaneously. In hypochlorous acid, the chlorine atom exists in a +1 oxidation state, which is unstable, causing it to shift to lower and higher oxidation states to reach more stable configurations. Consequently, the chlorine is reduced to a -1 oxidation state to form $\mathrm{HCl}$ and oxidized to a +5 oxidation state to form $\mathrm{HClO}_3$, resulting in the formation of hydrochloric acid and chloric acid.

Acidity in oxoacids is fundamentally linked to the oxidation state of the central atom, where a higher positive charge on that atom polarizes the oxygen-hydrogen bond more effectively, promoting proton dissociation. In the case of $\mathrm{HNO}_3$, the central nitrogen exists in a +5 oxidation state, whereas it is in a +3 state in $\mathrm{HNO}_2$. Because acid strength increases with the oxidation state of the central atom, $\mathrm{HNO}_3$ is a stronger acid than $\mathrm{HNO}_2$.

Aluminium chloride undergoes hydrolysis in water because it is a salt formed from a weak base and a strong acid, yielding an acidic solution containing aluminium hydroxide and hydrogen chloride, expressed as $\mathrm{AlCl}_3 + 3\mathrm{H}_2\mathrm{O} \rightarrow \mathrm{Al}(\mathrm{OH})_3 + 3\mathrm{HCl}$. As this aqueous solution is heated to dryness, the volatile hydrogen chloride escapes the mixture, while the remaining aluminium hydroxide thermally decomposes to leave behind aluminium oxide according to $2\mathrm{Al}(\mathrm{OH})_3 \xrightarrow{\Delta} \mathrm{Al}_2\mathrm{O}_3 + 3\mathrm{H}_2\mathrm{O}$.