P Block Elements – NEET UG ChemistryPractice Questions & PYQs

The readiness to form anions depends on an element's electron affinity and its proximity to achieving a stable, fully-filled valence shell configuration. Halogens possess seven valence electrons with an outer electronic configuration of $ns^2 np^5$, leaving them just one electron short of a stable octet. Because of this strong electronic preference, they readily accept an electron from metallic species to attain a noble gas configuration, allowing them to form halide anions with exceptional ease.

Elements belonging to the second period of the periodic table are strictly restricted to an octet of electrons in their valence shell due to the complete absence of vacant $d$-orbitals. While heavier group 15 congeners can utilize low-lying orbitals to expand their coordination numbers and form stable pentahalides, nitrogen lacks this capability and cannot accommodate five covalent bonds to form $\mathrm{NF}_5$.

The tendency to form anions is governed by high electronegativity and electron gain enthalpy, which are maximized in non-metals located toward the upper right of the periodic table. Elements such as $\mathrm{N}$, $\mathrm{O}$, and $\mathrm{F}$ possess compact atomic sizes and high nuclear charge, allowing them to readily attract and accept electrons to achieve a stable noble gas configuration. Metals like sodium, magnesium, gallium, and transition elements instead prefer to lose electrons and form cations. Therefore, the electronegative triad of nitrogen, oxygen, and fluorine exhibits the absolute strongest affinity for acquiring extra electrons.

Certain phenolic compounds in alkaline environments exhibit an extremely high affinity for gaseous oxygen due to rapid autoxidation pathways. An alkaline solution of pyrogallol rapidly strips oxygen from gas mixtures by chemically binding and consuming it, making it the standard laboratory reagent for gas analysis. Other drying agents or basic solutions lack this specific redox reactivity toward molecular oxygen.

Strong oxidizing agents like concentrated nitric acid readily attack complex carbohydrates by supplying nascent oxygen to break down the carbon backbone. When cane sugar undergoes this vigorous oxidation, the carbon-carbon skeleton and functional groups are systematically converted into a simple dicarboxylic acid. The balanced transformation can be represented as $\mathrm{C}_{12}\mathrm{H}_{22}\mathrm{O}_{11} + 18[\mathrm{O}] \rightarrow 6(\mathrm{COOH})_{2} + 5\mathrm{H}_{2}\mathrm{O}$, yielding oxalic acid as the primary organic product.

The basicity of an oxoacid is determined by the number of ionizable hydrogen atoms attached directly to oxygen atoms as hydroxyl groups. The chemical formula $\mathrm{H}_3 \mathrm{PO}_2$ corresponds to hypophosphorous acid, which structures its hydrogen atoms such that only one is bonded to oxygen in a $\mathrm{P}-\mathrm{OH}$ linkage while the remaining two bond directly to phosphorus. Because only hydrogen atoms bound to oxygen can be released as protons, the presence of a single hydroxyl group establishes a basicity of one.

An acid anhydride is an oxide derived from an acid by the complete removal of water. When two molecules of $\mathrm{HNO}_{3}$ undergo dehydration, a single molecule of $\mathrm{H}_{2}\mathrm{O}$ is eliminated, leaving behind the corresponding higher oxide of nitrogen. This direct stoichiometric relationship establishes $\mathrm{N}_{2}\mathrm{O}_{5}$ as the proper anhydride form of nitric acid.

Phosphorus trichloride undergoes rapid hydrolysis when exposed to moisture due to the high polarity of its covalent bonds. During this reaction, water molecules attack the central atom to replace chlorine substituents with hydroxyl groups. This transformation yields phosphorous acid alongside a mist of hydrogen chloride gas.

Salts derived from strong acids and weak bases hydrolyze to produce acidic solutions, which progressively lower the pH of agricultural soil with continuous application. Ammonium sulphate is formed from the neutralization of strong $\mathrm{H}_{2}\mathrm{SO}_{4}$ by weak $\mathrm{NH}_{4}\mathrm{OH}$. As plants absorb the ammonium cations for nutrition, the sulphate ions remain behind in the soil moisture, accumulating sulphuric acid and thereby increasing overall soil acidity over time.

Oxoacids of sulphur are characterized by specific structural units including $\mathrm{S}-\mathrm{OH}$ groups, terminal $\mathrm{S}=\mathrm{O}$ double bonds, $\mathrm{S}-\mathrm{O}-\mathrm{S}$ bridges, and $\mathrm{S}-\mathrm{O}-\mathrm{O}-\mathrm{S}$ peroxide linkages. Peroxodisulphuric acid ($\mathrm{H_2S_2O_8}$) contains a peroxide bond along with two $\mathrm{S}-\mathrm{OH}$ and four $\mathrm{S}=\mathrm{O}$ bonds, whereas standard sulphuric acid ($\mathrm{H_2SO_4}$) possesses two $\mathrm{S}-\mathrm{OH}$ and two $\mathrm{S}=\mathrm{O}$ bonds. Furthermore, pyrosulphuric acid ($\mathrm{H_2S_2O_7}$) features an $\mathrm{S}-\mathrm{O}-\mathrm{S}$ linkage with two $\mathrm{S}-\mathrm{OH}$ and four $\mathrm{S}=\mathrm{O}$ bonds, and sulphurous acid ($\mathrm{H_2SO_3}$) is structured with two $\mathrm{S}-\mathrm{OH}$ and one $\mathrm{S}=\mathrm{O}$ bond.

Deep-sea divers breathe a specialized mixture of oxygen and helium to prevent nitrogen narcosis and decompression sickness under high pressure. This protective benefit relies on helium having a very low solubility in biological fluids like blood, which contradicts the claim about high solubility. Consequently, the assertion regarding its use as a diluent is completely accurate, whereas the accompanying reason stating high solubility is entirely false.

Boiling points within a group of hydrides are heavily influenced by intermolecular forces rather than depending solely on molecular weight. Water ($\mathrm{H_2O}$) possesses an anomalously high boiling point due to extensive intermolecular hydrogen bonding networks that require substantial thermal energy to disrupt. As a result, the actual boiling point trend starts at the highest value for water and then increases from hydrogen sulfide ($\mathrm{H_2S}$) up to hydrogen telluride ($\mathrm{H_2Te}$), rendering the claimed simple monotonic order incorrect.

Interhalogen compounds generally exhibit higher chemical reactivity than diatomic halogens because of fundamental differences in their bond dissociation energies. Specifically, the covalent bond holding the iodine and chlorine atoms together ($\mathrm{I-Cl}$) is less stable and requires less energy to break than the homonuclear bond found in elemental iodine ($\mathrm{I-I}$). This lower bond energy facilitates much easier cleavage during chemical reactions, making the interhalogen species significantly more reactive than the pure halogen.

The industrial manufacture of nitric acid relies on the catalytic oxidation of ammonia as the primary foundational step of the Ostwald process. In this initial stage, gaseous ammonia reacts with atmospheric oxygen over a platinum-rhodium gauge catalyst under controlled thermal and pressure conditions. The specific chemical equation representing this catalytic conversion is given by:
$4\mathrm{NH}_3 + 5\mathrm{O}_2 \xrightarrow[500\text{ K}, \; 9\text{ bar}]{\mathrm{Pt/Rh}} 4\mathrm{NO} + 6\mathrm{H}_2\mathrm{O}$
This reaction efficiently yields nitric oxide, which is subsequently oxidized and absorbed in water to produce commercial-grade nitric acid.

The oxidizing strength of a halogen in aqueous solution is determined by the overall thermodynamic cycle of converting the diatomic gas into hydrated ions, encompassing bond dissociation, electron gain, and hydration. Fluorine exhibits an exceptionally high oxidizing capacity primarily because its diatomic molecules possess a weak bond that requires minimal energy to dissociate. Furthermore, the resulting fluoride ions release a massive amount of energy upon hydration owing to their small ionic radius and high charge density.

Noble gases are monoatomic species with completely filled valence shells, meaning they interact with one another solely through extremely weak London dispersion forces. Because these feeble intermolecular forces can be easily disrupted by minimal thermal energy, these elements characteristically exhibit very low melting and boiling points. This same electronic stability also accounts for their sparing solubility in water and their large positive values of electron gain enthalpy.

Acid strength in hydrohalic acids is fundamentally governed by the ease with which the hydrogen-halogen bond dissociates to release a proton. Moving down the group from fluorine to iodine, the atomic size increases substantially, leading to a longer bond length and a corresponding decrease in bond enthalpy. Consequently, this weaker bond makes it progressively easier for the molecule to donate a hydrogen ion, proving both statements regarding the increasing order of acidity to be true.

Oxoacids of sulphur possess diverse structural frameworks, among which specific peroxy derivatives contain a characteristic peroxide linkage connecting two structural units. Peroxodisulphuric acid, represented by $\mathrm{H}_2\mathrm{S}_2\mathrm{O}_8$, features an oxygen-oxygen single bond bridging two $-\mathrm{SO}_3\mathrm{H}$ groups together. This distinct peroxy bridge sets it apart from ordinary oxoacids such as sulphuric acid or pyrosulphuric acid, neither of which contains an $-\mathrm{O}-\mathrm{O}-$ unit.

Urea undergoes hydrolysis in the presence of water to yield ammonia, denoted as substance $\mathrm{B}$, which subsequently acts as a strong ligand. When this generated ammonia is introduced into an aqueous solution containing $\mathrm{Cu}^{2+}$ ions, it coordinates with the metal center to replace water molecules. This ligand exchange produces the deep blue tetraamminecopper(II) coordination complex with the formula $[\mathrm{Cu}(\mathrm{NH}_3)_4]^{2+}$.

The VSEPR theory determines the molecular geometry of xenon compounds by accounting for both bond pairs and lone pairs of electrons around the central xenon atom. Specifically, $XeF_4$ possesses six electron domains with two lone pairs leading to a square planar structure, whereas $XeF_6$ has seven domains with one lone pair resulting in a distorted octahedral geometry. Additionally, $XeOF_4$ adopts a square pyramidal shape due to five bond pairs and one lone pair, and $XeO_3$ forms a pyramidal structure with three bond pairs and one lone pair.