Some P Block Elements – NEET UG ChemistryPractice Questions & PYQs

Boric acid ($\mathrm{H}_3\mathrm{BO}_3$) functions as a Lewis acid because its central boron atom is electron-deficient and capable of accepting a lone pair of electrons. When dissolved in water, the molecule coordinates with a hydroxide ion provided by a water molecule to form the $[\mathrm{B}(\mathrm{OH})_4]^-$ complex. This reaction effectively releases the remaining proton from the water into the solution, demonstrating that the acidity arises from this specific interaction rather than simple dissociation.

Aluminum ions are limited to a maximum coordination number of six, meaning they readily accept available fluoride ligands to form stable octahedral complexes. In the presence of hydrogen fluoride and potassium fluoride, aluminum fluoride undergoes a complexation reaction that yields the soluble salt potassium hexafluoroaluminate, which is represented by the equation $\mathrm{AlF}_3 + 3\mathrm{KF} \xrightarrow{\mathrm{HF}} \mathrm{K}_3[\mathrm{AlF}_6]$ This chemical transformation facilitates the dissolution of $\mathrm{AlF}_3$ because the resulting complex anion is highly soluble within the solvent mixture.

Graphite is characterized by its distinct hexagonal layer structure, which arises from $sp^2$ hybridized carbon atoms arranged in sheets. Boron Nitride adopts an almost identical layered configuration, earning it the common designation of inorganic graphite because it is isoelectronic with carbon. This specific structural resemblance makes BN unique among the choices provided, as no other compound listed mimics that distinct planar arrangement.

The classification of silicate minerals is determined by how the fundamental $SiO_4^{4-}$ tetrahedra are linked to form distinct geometric structures. Feldspars are prominent minerals within this category, defined specifically as three-dimensional aluminosilicates where aluminum atoms substitute for silicon within the tetrahedral framework. Because this three-dimensional arrangement of aluminosilicates is the defining feature of feldspars, the claim that they are not aluminosilicates is factually incorrect.

Building high molecular mass silicone polymers relies on the ability of monomer units to link into continuous chains or networks, which requires each unit to have multiple reactive sites. These monomers need at least two hydrolysable chlorine atoms to allow for the stepwise propagation essential for long-chain growth. Since $Me_3SiCl$ features only a single chlorine atom, it cannot sustain chain formation and instead acts as a terminator, while species like $Me_2SiCl_2$ offer two reactive sites to create linear polymers, and $MeSiCl_3$ or $PhSiCl_3$ provide three reactive sites to build cross-linked structures.

Every silicate mineral is constructed from a fundamental tetrahedral building block known as the orthosilicate ion. Within this unit, a single silicon atom undergoes sp3 hybridization to establish covalent bonds with four oxygen atoms positioned at the vertices of a regular tetrahedron. This arrangement results in the chemical formula $\mathrm{SiO}_4^{4-}$, which functions as the universal structural repeat unit for the entire class of silicates.

Pure alumina ($Al_2O_3$) has a prohibitively high fusion temperature, making it difficult to electrolyze in its raw state. By incorporating cryolite ($Na_3AlF_6$) and fluorspar ($CaF_2$) into the electrolyte bath in a 20 : 40 : 2 ratio, the melting point is effectively lowered from 2000°C to 900°C. Furthermore, this deliberate mixture significantly increases the electrical conductivity of the molten solution, which is essential for efficient aluminium extraction.

Silicate structures are primarily classified by the number of oxygen atoms shared between the fundamental $[SiO_4]^{4-}$ tetrahedral building blocks. In the specific case of pyrosilicates, two tetrahedral units connect by sharing a single corner oxygen atom, which leads to the formation of the $[Si_2O_7]^{6-}$ ion. This simple linkage involves only one shared oxygen, distinguishing it from more complex structures where multiple corners are linked.

Amphoteric oxides are defined by their ability to behave as both acids and bases, reacting with either to form salts and water. Tin(IV) oxide, $\mathrm{SnO}_2$, demonstrates this dual reactivity, as shown by its interactions with acid and base in equations such as $\mathrm{SnO}_2 + 4\mathrm{HCl} \rightarrow \mathrm{SnCl}_2 + 2\mathrm{H}_2\mathrm{O}$ and $\mathrm{SnO}_2 + 2\mathrm{NaOH} \rightarrow \mathrm{Na}_2\mathrm{SnO}_3 + \mathrm{H}_2\mathrm{O}$. While these reactions establish the amphoteric nature of $\mathrm{SnO}_2$, the remaining substances behave differently, as $\mathrm{CaO}$ is basic, whereas $\mathrm{SiO}_2$ and $\mathrm{CO}_2$ are acidic in nature.

Aluminium displays amphoteric character, meaning it reacts with strong bases like sodium hydroxide, but this process does not stop at the formation of aluminium hydroxide. When aluminium metal reacts with excess $\mathrm{NaOH}$, it forms a soluble complex called sodium tetrahydroxoaluminate, expressed by the equation $2 \mathrm{Al} + 2 \mathrm{NaOH} + 6 \mathrm{H}_2\mathrm{O} \longrightarrow 2 \mathrm{NaAl}(\mathrm{OH})_4 + 3 \mathrm{H}_2$. Since this reaction results in a soluble aluminate rather than a hydroxide precipitate, the claim that it produces $\mathrm{Al}(\mathrm{OH})_3$ is chemically inaccurate. Meanwhile, the other statements are correct: sodium dissolves in liquid ammonia to yield a blue solution of solvated electrons, sodium hydroxide corrodes glass by forming sodium silicate, and heating sodium bicarbonate releases carbon dioxide to leave behind sodium carbonate.

A Lewis acid is defined by the ability to accept a lone pair of electrons, typically to fulfill an octet or form a coordinate bond. Molecules such as $NH_3$ and $H_2O$ possess lone pairs on their central atoms, functioning instead as electron-donating Lewis bases, whereas $CH_4$ is a stable, neutral molecule that does not engage in these electron pair interactions. Conversely, $B_2H_6$ contains electron-deficient banana bonds and vacant orbitals, which gives it the distinct capacity to act as an electron pair acceptor, qualifying it as a Lewis acid.

The Lewis acidity of boron trihalides is primarily determined by the extent of back-donation from the halogen's filled $p$-orbitals into the empty $2p$-orbital of the boron atom. In $\mathrm{BF}_{3}$, the small fluorine atom allows for excellent $2p-2p$ orbital overlap, facilitating strong $p\pi-p\pi$ back-bonding that effectively minimizes boron's electron deficiency. As the halogen size increases and the energy match with boron's orbitals worsens, this stabilizing back-donation diminishes, reaching its minimum in $\mathrm{BBr}_{3}$. Because $\mathrm{BBr}_{3}$ experiences the least internal stabilization, the boron center remains the most electrophilic, establishing the overall Lewis acidity trend as $\mathrm{BBr}_{3} > \mathrm{BCl}_{3} > \mathrm{BF}_{3}$.

The Group 13 elements possess a valence shell configuration of $ns^2 np^1$, but as we move down the group, the $ns^2$ electrons are increasingly reluctant to participate in bonding due to poor shielding by inner orbitals, a phenomenon known as the inert pair effect. This trend makes the loss of only the single $np^1$ electron more energetically favorable compared to the loss of all three valence electrons. As a result, the stability of the +1 oxidation state increases down the group, following the sequence $\mathrm{Al} < \mathrm{Ga} < \mathrm{In} < \mathrm{Tl}$.

The formation of straight-chain polymers from chlorosilanes depends on the functionality of the monomer, which is determined by the number of chlorine atoms attached to the silicon center. Because $(CH_3)_2SiCl_2$ contains two chlorine atoms, its hydrolysis results in a dihydroxy silanol, which acts as a bifunctional building block capable of polymerizing into a long, linear chain through condensation. Monomers with only one chlorine atom terminate chains, while those with three chlorine atoms create complex cross-linked networks, making the bifunctional nature of $(CH_3)_2SiCl_2$ the necessary requirement for linear growth.

The stability of the +1 oxidation state in Group 13 elements is governed by the inert pair effect, where the valence $ns^2$ electrons become progressively more reluctant to participate in bonding as the principal quantum number increases. Because this reluctance to bond is most pronounced in heavier elements, the +1 state becomes thermodynamically more favorable relative to the +3 state as we descend the group. This behavior results in a clear trend of increasing stability for the +1 oxidation state in the order $Al^{+} < Ga^{+} < In^{+} < Tl^{+}$.

The fundamental building block of all silicate minerals is the $SiO_4^{4-}$ tetrahedron, and the specific geometry of the structure depends on how these units link together. In chain silicates, each individual tetrahedron shares two of its oxygen atoms with neighboring units to form continuous, linear arrangements. This specific connectivity results in a repeating anionic unit with the general formula $(SiO_3)^{2n-}_n$.

The inert pair effect describes the reluctance of the outermost $ns^2$ electrons to participate in bonding, a phenomenon that becomes more pronounced as we descend Group 14. Since this effect strengthens as atomic size increases, the stability of the higher $+4$ oxidation state decreases down the group, while the stability of the lower $+2$ oxidation state increases. Consequently, the $+2$ oxidation state is the most characteristic for lead, while the $+4$ oxidation state is the most characteristic for tin.

Electronegativity is fundamentally linked to the percentage of s-character in a hybrid orbital because s-orbitals are spatially closer to the positively charged nucleus, allowing them to exert a stronger pull on bonding electrons. Because the $sp$ hybrid orbital has the highest s-character, followed by $sp^2$ and then $sp^3$, the capacity of these orbitals to attract electrons decreases in that same sequence. This relationship dictates that the electronegativity of carbon hybrid orbitals follows the order $sp > sp^2 > sp^3$.

The acidity or basicity of an oxide is fundamentally linked to the metallic character of the element, where non-metallic oxides are typically acidic, and metallic oxides become increasingly basic as the element's metallic nature strengthens. Selenium, a non-metal, produces the acidic oxide ${SeO}_{2}$, while both ${Al}_{2}{O}_{3}$ and ${Sb}_{2}{O}_{3}$ exhibit amphoteric behavior, meaning they can react with both acids and bases. Because bismuth sits at the bottom of its group and possesses the most significant metallic character among these choices, its oxide, ${Bi}_{2}{O}_{3}$, is the only one that functions as a basic oxide.

Zeolites are aluminosilicate materials characterized by a stable, three-dimensional framework built exclusively from $SiO_4$ and $AlO_4$ tetrahedra. While the replacement of silicon with aluminum in these units generates a net negative charge that allows these materials to function as cation exchangers, the claim that $AlO_6^{9-}$ ions exist within this structure is false because the framework is defined by tetrahedral coordination. This distinct, open architecture creates cavities throughout the solid, which is precisely what enables zeolites to act as molecular sieves capable of taking up water and other small molecules.