Some P Block Elements – NEET UG ChemistryPractice Questions & PYQs

The name "lead pencil" is a historical misnomer because these writing instruments contain absolutely no metallic lead. The core is actually crafted from a mixture of graphite and clay, which is the substance responsible for leaving a mark on paper. Because the composition is entirely restricted to these non-lead materials, the percentage of lead in a lead pencil is zero.

The ability of a material to conduct electricity relies on the presence of mobile charge carriers, typically free electrons. Unlike aluminium, graphite, and silicon, which act as conductors because they possess these mobile carriers, diamond maintains a rigid crystalline structure where all valence electrons are tightly locked in strong covalent bonds. Because these electrons are unavailable to move through the lattice, diamond cannot facilitate the flow of current and functions as an insulator.

Because carbon atoms in diamond adopt $sp^3$ hybridization, they are able to share valence electrons with four distinct neighboring carbon atoms to achieve a stable configuration. These atoms are geometrically oriented toward the vertices of a regular tetrahedron, and this mechanism of sharing electron pairs defines the interaction as a covalent bond. This rigid, interconnected network serves as the fundamental basis for the material's structure.

When a chemical substance possesses the inherent capability to organize itself into two or more distinct crystalline arrangements while maintaining the same chemical composition, this phenomenon is identified as polymorphism. This structural variability arises because the same material can adopt different spatial patterns based on varying environmental conditions such as temperature and pressure. Unlike isomorphism, which involves different substances sharing the same crystalline shape, or isomerism, which relates to molecular connectivity, polymorphism specifically describes the diverse crystalline manifestations of a single substance.

Diborane ($B_2H_6$) features each boron atom bonded to four surrounding atoms, specifically two terminal hydrogen atoms and two bridging hydrogen atoms. To accommodate this coordination number of four, each boron atom must undergo $sp^3$ hybridization to generate four equivalent hybrid orbitals. In this configuration, three of these orbitals contain a single electron each to participate in terminal bonding, while the fourth orbital remains empty, allowing it to interact with a bridging hydrogen atom and the other boron atom to form the characteristic three-centered, two-electron bonds.

Electrical conduction relies on the presence of free, mobile charge carriers, such as delocalized electrons, within a substance. Potassium and sodium are alkali metals that conduct electricity effectively because their metallic bonding creates a sea of free electrons. Graphite also conducts electricity despite being a non-metal, as its layered structure allows one valence electron from each carbon atom to move freely across the planes. In contrast, diamond consists of a rigid, three-dimensional tetrahedral lattice where all four valence electrons of every carbon atom are locked into strong covalent bonds, leaving no free electrons to facilitate the flow of current.

Potash alum is a double salt defined by the chemical formula $\mathrm{K}_2\mathrm{SO}_4 \cdot \mathrm{Al}_2(\mathrm{SO}_4)_3 \cdot 24\mathrm{H}_2\mathrm{O}$. Within this structure, there are two $\mathrm{K}^{+}$ ions sourced directly from the potassium sulfate component. The sulfate ions are distributed across the formula, with one derived from $\mathrm{K}_2\mathrm{SO}_4$ and three from $\mathrm{Al}_2(\mathrm{SO}_4)_3$, totaling four $\mathrm{SO}_4^{2-}$ ions. This results in an ion count ratio of $2:4$, which simplifies to $1:2$.

In the structure of graphite, each carbon atom undergoes $sp^2$ hybridization to form covalent single bonds with three neighboring carbon atoms. The fourth valence electron from each carbon contributes to the formation of $\pi$ bonds that are not fixed to any single bond location. Because of this, the electrons are delocalized and spread out across the layered structure, which maintains a separation of 340 pm between sheets. This delocalization provides the conductive properties associated with the graphite lattice.

Aluminum trichloride acts as an electron-deficient species in its monomeric form, leaving the central aluminum atom with an incomplete octet. By utilizing vacant $3d$-orbitals, the aluminum atom can accept electron pairs from chloride ions to establish coordinate covalent bonds, which increases its coordination number and effectively completes its octet. This structural ability to expand its covalency is precisely what drives $\text{AlCl}_3$ to pair up and form a stable dimer.

Molecular geometry is fundamentally governed by the valence shell electron pair repulsion (VSEPR) theory, which dictates that lone pairs of electrons exert significant repulsive forces on surrounding bond pairs. In $\text{BCl}_3$, the central boron atom is surrounded by three bonding pairs with no remaining lone electrons, leading to a symmetrical, flat trigonal planar structure. In contrast, the nitrogen atom in $\text{NCl}_3$ possesses a lone pair of electrons that actively repels the three $\text{N}-\text{Cl}$ bonds, distorting the molecular framework into a non-planar, pyramidal geometry.

Redox chemistry centers on the change in oxidation states of elements during a chemical reaction, where a decrease in oxidation number signifies reduction and an increase signifies oxidation. When sodium chloride is heated with concentrated $\mathrm{H}_2\mathrm{SO}_4$ and $\mathrm{MnO}_2$, the reaction yields $\mathrm{Cl}_2$ and $\mathrm{MnSO}_4$, causing manganese to transition from a +4 oxidation state in $\mathrm{MnO}_2$ to a +2 state in $\mathrm{MnSO}_4$. This decrease in oxidation number confirms that manganese undergoes reduction rather than oxidation, rendering the first statement incorrect. Conversely, in the reaction involving sodium iodide, the same reagents drive the oxidation of iodide to iodine while manganese again shifts from a +4 to a +2 oxidation state, verifying that the second statement correctly describes the reduction of manganese.

Nitrogen, as a period 2 element, is restricted to s and p orbitals and completely lacks d-orbitals in its valence shell, making it impossible for the atom to engage in bonding schemes that rely on d-orbital participation, such as the proposed $d\pi-p\pi$ interaction with oxygen. While its small size and electronic configuration enable nitrogen to form strong $p\pi-p\pi$ multiple bonds with itself, this lack of vacant orbitals fundamentally precludes the formation of $d\pi-p\pi$ bonds. In contrast, heavier elements in the group such as phosphorus, arsenic, and antimony possess vacant d-orbitals that facilitate catenation and allow them to form $d\pi-d\pi$ bonds with transition metals, rendering the statement regarding nitrogen's ability to form $d\pi-p\pi$ bonds with oxygen incorrect.

A Lewis acid is formally defined as a chemical species capable of accepting a lone pair of electrons to complete its valence shell. Compounds of boron, like $BF_3$, possess a central atom that is surrounded by only six valence electrons, leaving it with an incomplete octet. This inherent electron deficiency makes boron compounds highly reactive toward electron donors, as they seek to fill this vacancy and stabilize their electronic structure, which serves as the fundamental basis for their classification as Lewis acids.

When ions are isoelectronic, they possess the same total number of electrons, making the nuclear charge the primary factor governing their size. Moving from ${Na}^{+}$ to ${Si}^{4+}$, the increasing number of protons exerts a stronger electrostatic pull on the electron cloud, drawing it closer to the nucleus. This increased effective nuclear charge causes the ionic radius to shrink as the positive charge increases, establishing the size order as ${Na}^{+} > {Mg}^{2+} > {Al}^{3+} > {Si}^{4+}$.

The coordination number of an atom is primarily determined by the availability of orbitals within its valence shell to accommodate bonding pairs. Carbon is restricted to a maximum coordination number of 4 because it lacks $d$-orbitals, which limits its bonding capacity to the $2s$ and $2p$ orbitals. In contrast, silicon belongs to the third period and possesses vacant low-lying $d$-orbitals, which allow it to expand its valence shell and accommodate a higher coordination number of 6.

Glass is scientifically classified as a supercooled liquid because its internal molecular structure lacks the ordered, repeating pattern required to form a true crystalline solid. During the cooling process, the atoms are frozen into a disordered state rather than organizing into a regular lattice, which results in the material's distinct amorphous properties. This specific lack of structure, despite the material's rigid appearance, defines it as a liquid that has been cooled without undergoing a standard crystallization phase.

Boric acid functions as a weak monobasic acid rather than a strong tribasic one because it does not simply dissociate to donate three protons in solution. Its acidic behavior stems from its role as a Lewis acid, where it accepts a hydroxyl ion from water to produce the $[{B}( {OH})_{4}]^{-}$ anion while releasing a single proton, expressed as ${B}( {OH})_{3} + {H}_{2} {O} \rightarrow [{B}( {OH})_{4}]^{-} + {H}^{+}$. Consequently, labeling it as a strong tribasic acid is incorrect, as the compound is neither a strong electrolyte nor capable of donating three hydrogen ions.

Silicon tetrachloride ($SiCl_4$) is a chemical compound characterized by its rapid reactivity with water vapor present in the atmosphere. When exposed to moist air, it undergoes hydrolysis to produce thick, opaque white fumes composed of finely divided particles. These dense clouds effectively scatter light and obstruct visibility, making $SiCl_4$ a practical substance for generating a tactical smoke screen in warfare.

The maximum covalency of an element is determined by the total number of valence orbitals available to participate in bonding, which allows atoms to potentially exceed their group valency when empty orbitals are accessible. Boron, having three valence electrons and an empty 2p orbital, can accept a lone pair of electrons to form a maximum of four bonds, such as in $\mathrm{BF_4^-}$. Aluminium, on the other hand, contains nine valence orbitals comprising the 3s, three 3p, and five 3d orbitals, meaning its bonding capacity is not limited to five. While Beryllium contains four valence orbitals within its 2s and three 2p subshells, its maximum covalency is restricted to two, whereas Magnesium can expand its octet by utilizing 3d orbitals to achieve a covalency of six, as observed in $\mathrm{[Mg(NH_3)_6]^{2+}}$.

Group 15 elements consistently react with hydrogen to produce hydrides with the general formula $\mathrm{EH}_3$, which confirms that nitrogen forms ammonia $(\mathrm{NH}_3)$ and arsenic forms arsine $(\mathrm{AsH}_3)$ as described in the first statement. These same elements possess the ability to create two distinct series of oxides, specifically $\mathrm{E}_2\mathrm{O}_3$ and $\mathrm{E}_2\mathrm{O}_5$; since antimony is capable of producing antimony pentoxide $(\mathrm{Sb}_2\mathrm{O}_5)$, the assertion that it cannot do so is inaccurate.