P Block Elements – NEET UG ChemistryPractice Questions & PYQs

Metal phosphides react readily with water to liberate phosphine gas, which is widely recognized for its characteristic pungent odor resembling rotting fish. When this colorless gas, $\mathrm{PH_3}$, is bubbled through a copper sulfate solution, it undergoes a precipitation reaction to form copper phosphide along with sulfuric acid. Tracing these reaction pathways backward reveals that the initial reactant yielding phosphine upon hydrolysis must be calcium phosphide. Thus, the compound $\mathrm{X}$ is $\mathrm{Ca_3P_2}$, which acts as the primary precursor for generating phosphine in such synthetic pathways.

Thermal stability within group sixteen hydrides is dictated by the bond dissociation enthalpy of the $\mathrm{H}-\mathrm{E}$ bond, which naturally diminishes as the central atom increases in size down the group. Because larger elements form longer and progressively weaker covalent bonds with hydrogen, the molecules require significantly less thermal energy to break apart. Thus, the thermal stability decreases sequentially from $\mathrm{H}_2\mathrm{O}$ down to $\mathrm{H}_2\mathrm{Po}$.

The reducing strength of phosphorus oxoacids depends directly on the number of hydrogen atoms directly bonded to the phosphorus atom, as these specific linkages are easily oxidized. Inspecting the molecular structures shows that hypophosphorous acid possesses two such hydrogen atoms linked directly to the central atom, whereas other acids in this category contain fewer. Because it features the maximum count of these labile bonds among the choices, it readily donates electrons to reduce other chemical species. Consequently, $\mathrm{H}_3 \mathrm{PO}_2$ exhibits the most powerful reducing capabilities in this series.

Oleum is formed by dissolving sulfur trioxide gas directly into concentrated sulfuric acid, resulting in a dense, fuming liquid consisting largely of pyrosulfuric acid. This compound acts as an important industrial intermediate and sulfonating agent, possessing a condensed dimeric structure of sulfate units joined by an oxygen bridge. Balancing the proportions of sulfur trioxide and sulfuric acid yields the specific molecular formula containing two sulfur atoms and seven oxygen atoms. Therefore, the chemical representation corresponding to this fuming mixture is $\mathrm{H}_2\mathrm{S}_2\mathrm{O}_7$.

Pure nitrogen is obtained in a highly pure form through the thermal decomposition of sodium azide or barium azide. The Haber process serves as the industrial method for synthesizing ammonia, whereas the Contact process is used for the large-scale production of sulphuric acid. Finally, Deacon's process is employed for the manufacture of chlorine gas via the catalytic oxidation of hydrogen chloride.

The molecular geometry of $PCl_5$ is trigonal bipyramidal, which forces the five chlorine atoms into two distinct electronic environments with varying bond lengths and angles. Specifically, the three equatorial P-Cl bonds lie in a plane making angles of 120° with each other, whereas the two axial P-Cl bonds lie perpendicular to this plane at an angle of 180°. Because the axial bonds experience greater repulsion from the equatorial bond pairs, they are significantly longer (240 pm) and weaker than the equatorial bonds (202 pm). This inherent bond weakness and length make the molecule quite chemically reactive rather than stable and inert.

Halogens exhibit diverse chemical behaviors across their group, yet almost all of them can display positive oxidation states in addition to their characteristic negative states when combining with more electronegative elements. While fluorine is traditionally known for being exclusively negative due to its extreme electronegativity, it actually adopts a positive oxidation state of plus one in compounds like hypofluorous acid. This unique exception means that fluorine, just like the heavier halogens, is capable of exhibiting positive oxidation states under specific bonding conditions with oxygen. Consequently, the claim that every halogen except fluorine displays positive oxidation states is inaccurate.

VSEPR theory determines the molecular geometry of interhalogen compounds based on the total number of bond pairs and lone pairs around the central halogen atom. A diatomic interhalogen molecule of type $XX'$ naturally adopts a linear geometry. When five electron pairs are present with two lone pairs as in $XX'_3$, the molecule bends into a T-shape, whereas six electron pairs with one lone pair in $XX'_5$ result in a square pyramidal structure. Finally, seven electron pairs around the central atom in $XX'_7$ force a pentagonal bipyramidal arrangement to minimize electronic repulsion.

The structural framework of polythionate and thiosulfate anions reveals the presence of direct sulfur-sulfur covalent linkages. Tetrathionate ions contain a central chain of four sulfur atoms linked together, meaning an explicit sulfur-sulfur bond is present within the backbone. Similarly, thiosulfate ions feature a central sulfur atom bonded directly to another sulfur atom via a single covalent bond alongside three terminal oxygen atoms. Other oxoanions like disulfate or peroxodisulfate lack direct sulfur-sulfur bonds because their sulfur centers are bridged exclusively through oxygen atoms.

Concentrated nitric acid acts as a powerful oxidizing agent that readily oxidizes transition metals while being reduced itself to nitrogen dioxide gas. When copper metal reacts with this concentrated acid, the copper is converted into $\mathrm{Cu}(\mathrm{NO}_3)_2$ alongside the evolution of dense brown fumes of $\mathrm{NO}_2$ and water vapor. In contrast, dilute nitric acid yields nitric oxide instead, highlighting the concentration-dependent oxidizing capability of the acid.

VSEPR theory governs the three-dimensional architecture of xenon compounds by accounting for both bonding electron pairs and steric lone pairs on the central xenon atom. Xenon hexafluoride possesses seven electron pairs including a lone pair, which distorts its octahedral geometry into a distorted octahedral shape. Xenon trioxide features three bond pairs and one lone pair leading to a pyramidal structure, while xenon oxotetrafluoride accommodates six electron pairs with a lone pair to form a square pyramidal geometry. Finally, xenon tetrafluoride adopts a square planar configuration because its six electron pairs arrange with two lone pairs positioned at opposite vertical poles to minimize repulsion.

Bond dissociation enthalpy in halogens generally decreases down the group due to increasing atomic size, except for fluorine, which has an abnormally low value caused by severe electronic repulsion between lone pairs in its small valence shell. Chlorine possesses an optimal bond length and effective orbital overlap, giving it the highest bond dissociation energy among the halogens. As atomic radius increases further down to bromine and iodine, the bond length elongates significantly and the overlap becomes diffuse, causing the dissociation enthalpy to drop steadily. Consequently, the overall trend peaks at chlorine and follows a downward progression through bromine, fluorine, and finally iodine at the lowest extreme.

The basicity of an oxoacid of phosphorus is determined strictly by the number of hydrogen atoms directly attached to oxygen atoms, which are ionizable as protons. Phosphinic acid contains only one hydroxyl group bonded to phosphorus, making it a monoprotic acid capable of releasing a single hydrogen ion. In contrast, phosphonic acid features two hydroxyl groups bonded to the central phosphorus atom, granting it a diprotic nature. Hydrogen atoms bonded directly to the phosphorus atom are non-ionizable and do not contribute to the overall basicity of the molecule.

Nomenclature rules for binary chemical compounds dictate that the more electronegative element is always designated as the suffix root, while the less electronegative element acts as the central prefix. Since fluorine has a higher electronegativity than oxygen, any binary compound combining them must be named as a fluoride of oxygen rather than an oxide of fluorine. Consequently, oxygen difluoride correctly categorizes $\mathrm{OF}_2$, making the designation of it as an oxide scientifically inaccurate.

Sulphur dioxide ($\mathrm{SO_2}$) and nitrogen dioxide ($\mathrm{NO_2}$) share several chemical traits, including water solubility to form acidic solutions, functioning as reducing agents, and contributing to the formation of acid rain. However, sulphur dioxide is widely employed as a preservative in food and beverages due to its antimicrobial and antioxidant properties. Nitrogen dioxide possesses toxic characteristics and strong oxidizing traits at higher concentrations that preclude its use in food preservation.

Intermolecular forces heavily dictate the physical properties of covalent hydrides, with exceptionally strong dipole interactions causing major anomalies in boiling points. Fluorine possesses an extremely high electronegativity and a very small atomic radius, which allows hydrogen fluoride molecules to associate through robust hydrogen bonds. Overcoming these extensive intermolecular networks requires a significantly greater amount of thermal energy than merely overcoming the weaker van der Waals forces found in other hydrogen halides.

The reducing power of phosphorus oxyacids is directly related to the presence of hydrogen atoms directly attached to the phosphorus atom rather than through oxygen. In hypophosphorous acid, or $\mathrm{H}_3\mathrm{PO}_2$, the molecular structure contains two covalent $\mathrm{P}-\mathrm{H}$ bonds and only one ionizable $-\mathrm{OH}$ group. These directly bonded hydrogen atoms are easily oxidized, imparting strong reducing characteristics to the molecule.

Adding an electron to a singly charged negative ion introduces significant electrostatic repulsion within the compact electron cloud. Although acquiring a full octet gives the resulting doubly charged ion a stable electronic configuration matching neon, the energy penalty required to overcome inter-electronic repulsion is exceptionally high. This overwhelming repulsive force vastly outweighs any thermodynamic stability gained from achieving noble gas configuration, making the second electron addition strongly endothermic.

Acid strength down a group in the periodic table is primarily governed by the decreasing bond dissociation enthalpy resulting from an increase in atomic size. As we move from $\mathrm{H_2S}$ to $\mathrm{H_2Te}$, the bond length increases significantly, making it progressively easier for the molecule to release hydrogen ions in aqueous solution. Consequently, the ease of proton donation escalates down the group, establishing $\mathrm{H_2Te}$ as the strongest acid and $\mathrm{H_2S}$ as the weakest among the three hydrides.

The acidity of oxoacids is determined by the central atom's electronegativity, its oxidation state, and the stability of the resulting conjugate base through charge delocalization. Perchloric acid ($\mathrm{HClO_4}$) features chlorine in its maximum oxidation state of plus seven alongside extensive resonance stabilization across four oxygen atoms in its conjugate base. These combined structural factors allow the molecule to release a proton far more readily than the sulphur or lower chlorine oxoacids listed.