p-Block Elements — Practice Questions with Answers

Boron in $BF_3$ has only 6 electrons around it (an incomplete octet), leaving an empty 2p orbital available on the boron atom.

This empty orbital makes $BF_3$ highly capable of accepting an electron pair from an electron-rich species (a Lewis base) to complete boron's octet — this strong tendency to accept an electron pair is precisely why $BF_3$ acts as a strong Lewis acid.

The inert pair effect refers to the reluctance of the outermost $ns^2$ electron pair to participate in bonding, becoming increasingly pronounced for heavier elements down a group, due to poor shielding by intervening d and f electrons and relativistic effects, resulting in these electrons being held more tightly closer to the nucleus.

In Group 14, this effect means heavier elements like tin (Sn) and especially lead (Pb) increasingly favour the +2 oxidation state (where the $ns^2$ pair remains unused/inert) over the +4 state (which would require using all four valence electrons) — this is why $PbO$ (Pb²⁺) is notably more stable than $PbO_2$ (Pb⁴⁺), contrary to the group trend seen with lighter members like carbon, which strongly favours +4.

Diamond has each carbon atom covalently bonded ($sp^3$ hybridised) to four other carbon atoms in a rigid, extended three-dimensional tetrahedral network — this continuous, strong covalent bonding throughout the entire crystal structure makes diamond exceptionally hard.

Graphite has each carbon atom ($sp^2$ hybridised) bonded to only three other carbons within flat, hexagonal layers — while bonding within each layer is strong, the separate layers are held together only by weak van der Waals forces, allowing layers to easily slide over one another, giving graphite its characteristic softness and lubricating, slippery properties.

Nitrogen's small atomic size allows effective sideways overlap of its p-orbitals, enabling it to form a very strong triple bond ($N\equiv N$) with another nitrogen atom — this exceptionally strong triple bond makes the diatomic $N_2$ molecule highly stable, favouring this form over larger polyatomic structures.

Phosphorus, being significantly larger, cannot achieve effective p-orbital overlap for comparable, strong pi bonding — instead, phosphorus atoms form single bonds with several neighbouring phosphorus atoms, resulting in the tetrahedral $P_4$ molecular structure (in white phosphorus) as its characteristically stable form, rather than a diatomic species.

Ozone has a bent, three-atom structure with a resonance-stabilised structure involving one O-O single bond and one O=O double bond character (a bond order of 1.5 for each O-O bond, intermediate between single and double).

The decomposition of ozone into ordinary, more stable diatomic oxygen ($2O_3 \rightarrow 3O_2$) is thermodynamically favourable (releases energy), since $O_2$'s stronger, full double bond (bond order 2) represents a lower-energy, more stable bonding arrangement compared to ozone's comparatively weaker, resonance-delocalised bonds — this is why ozone is inherently less stable than ordinary oxygen, and tends to spontaneously decompose over time.

In hot, concentrated sulfuric acid, sulfur is present in its highest possible oxidation state (+6). Under these conditions, sulfuric acid can readily accept electrons from a suitable reducing agent, with sulfur being reduced to a lower oxidation state (commonly to $SO_2$, sulfur dioxide, +4 oxidation state).

This ready ability of sulfur to be reduced (accepting electrons) from its high +6 oxidation state, particularly under hot, concentrated conditions, is precisely why concentrated sulfuric acid acts as a powerful oxidising agent in such reactions (e.g., its reaction with copper metal, which it can oxidise despite copper's low reactivity with dilute acids).

Fluorine's very small atomic size means the two fluorine atoms in an $F_2$ molecule are held very close together, causing the several lone pairs of electrons on each fluorine atom to experience significant mutual repulsion with the lone pairs on the adjacent fluorine atom.

This substantial inter-electronic (lone pair-lone pair) repulsion in the compact $F_2$ molecule significantly weakens the F-F bond, resulting in a lower bond dissociation enthalpy than might otherwise be expected — chlorine, being a larger atom, experiences comparatively less of this specific lone pair repulsion effect in $Cl_2$, resulting in a stronger, higher bond dissociation enthalpy for Cl-Cl despite chlorine's lower electronegativity.

Noble gas atoms are monatomic and non-polar, held together in the liquid/solid state only by weak, temporary van der Waals (London dispersion) forces — since these forces are considerably weaker than typical covalent or ionic bonding, very little thermal energy is needed to overcome them, resulting in very low boiling points.

Down the group, atomic size (and hence the number of electrons and overall electron cloud polarisability) increases — larger, more polarisable electron clouds allow for stronger (though still relatively weak) instantaneous dipole-induced dipole interactions between atoms, resulting in progressively stronger intermolecular forces and hence increasing boiling points down the noble gas group.

Catenation is the ability of atoms of the same element to form extended chains or rings by bonding to each other. Carbon exhibits this property to an extraordinary extent, forming the vast, diverse array of organic compounds that constitutes organic chemistry.

Carbon's small size allows it to form exceptionally strong, stable C-C single bonds (as well as C=C and C≡C multiple bonds), supporting long, stable chains. Silicon, being significantly larger, forms comparatively weaker Si-Si bonds (due to poorer orbital overlap at greater distance) — these weaker bonds are less capable of supporting extensive, stable chain formation, which is why silicon shows far more limited catenation compared to carbon's extraordinarily extensive catenation capability.

Xenon is a relatively large noble gas atom, meaning its outermost electrons are held comparatively less tightly by the nucleus (larger atomic radius, more effective inner-shell shielding) than the smaller noble gases (like helium, neon, argon) — this gives xenon a comparatively lower (though still high) ionisation enthalpy among noble gases.

When combined with highly electronegative, small elements (particularly fluorine, and to a lesser extent oxygen) capable of forming strong bonds even with a reluctant partner, xenon's comparatively lower ionisation enthalpy allows these specific elements to successfully induce limited chemical bonding — forming a small number of genuinely stable compounds (like $XeF_2$, $XeF_4$, $XeF_6$) — a capability essentially unavailable to the smaller, higher-ionisation-enthalpy noble gases positioned above xenon in the group.

Diborane has the structure with two boron atoms bridged by two hydrogen atoms, along with two additional terminal hydrogen atoms on each boron (4 terminal + 2 bridging = 6 total hydrogens).

Since boron has only 3 valence electrons and forms only 4 terminal B-H bonds (2 per boron) using conventional 2-centre-2-electron bonds, there aren't enough electrons available for conventional bonding to also connect the two bridging hydrogens using standard bonds.

Instead, the two bridging B-H-B interactions are described as "banana bonds" or 3-centre-2-electron bonds — each bridging bond involves only 2 electrons shared collectively across 3 atomic centres (one hydrogen and two borons), rather than 2 electrons per single bond as in conventional bonding — this electron-deficient bonding arrangement is a distinctive, characteristic feature of diborane and related boron hydride compounds.

Haber process reaction: $N_2(g)+3H_2(g) \rightleftharpoons 2NH_3(g)$, $\Delta H<0$ (exothermic)

Conditions used: High pressure (typically 150-200 atm), moderate temperature (approximately 450°C), and an iron catalyst (with promoters like $Mo$).

Justification: High pressure favours the forward reaction (fewer moles of gas on the product side, per Le Chatelier's principle), maximising ammonia yield. Although lower temperature would thermodynamically favour the exothermic forward reaction even more (per Le Chatelier), a moderate (rather than very low) temperature is used as a practical compromise to ensure a reasonably fast reaction rate — the iron catalyst further speeds up the approach to equilibrium without affecting the equilibrium position itself, making the overall process economically viable within a reasonable timeframe.

Step 1: Sulfur is burned in air to produce sulfur dioxide: $S+O_2 \rightarrow SO_2$

Step 2: Sulfur dioxide is catalytically oxidised to sulfur trioxide, using a vanadium pentoxide ($V_2O_5$) catalyst at approximately 450°C: $2SO_2+O_2 \xrightarrow{V_2O_5} 2SO_3$

Step 3: Sulfur trioxide is absorbed in concentrated sulfuric acid (not directly in water, to avoid forming a difficult-to-control acid mist) to form oleum ($H_2S_2O_7$): $SO_3+H_2SO_4 \rightarrow H_2S_2O_7$

Step 4: Oleum is then carefully diluted with the appropriate amount of water to produce sulfuric acid of the desired concentration: $H_2S_2O_7+H_2O \rightarrow 2H_2SO_4$

Xenon in $XeF_4$ has 4 bond pairs (with 4 fluorine atoms) and 2 lone pairs of electrons — a total of 6 electron pairs around the central xenon atom.

Hybridisation: $sp^3d^2$ (mixing 1 s, 3 p, and 2 d orbitals to accommodate 6 electron pairs).

Underlying electron pair geometry: Octahedral (for 6 total electron pairs). To minimise repulsion, the 2 lone pairs occupy positions directly opposite each other (at 180°, the axial positions), leaving the 4 bond pairs (to fluorine) in the remaining equatorial plane.

Molecular shape (considering only bonded atoms): Square planar, with all four Xe-F bonds lying in the same plane, at 90° to each other.

Acidic strength (tendency to release $H^+$ in aqueous solution) of hydrogen halides increases in the order: HF < HCl < HBr < HI

This trend is explained by decreasing H-X bond dissociation enthalpy down the group — as halogen atomic size increases down the group (F→Cl→Br→I), the H-X bond becomes progressively longer and weaker (bond dissociation enthalpy decreases), making it progressively easier for the bond to break and release $H^+$ ions in solution.

Although HF might be expected to be a very strong acid based on fluorine's high electronegativity, its unusually strong, short H-F bond (along with strong hydrogen bonding effects) actually makes HF a comparatively weak acid — an important exception that must be explained via bond strength considerations, rather than simple electronegativity arguments alone.

When chlorine gas is dissolved in water, it undergoes a disproportionation reaction: $Cl_2+H_2O \rightleftharpoons HCl+HOCl$ (hypochlorous acid)

The hypochlorous acid (HOCl) formed is a powerful, effective germicidal (disinfecting) agent — it readily releases nascent oxygen/oxidising species that effectively kill harmful bacteria and other disease-causing microorganisms present in the water.

This germicidal action of the hypochlorous acid byproduct is precisely why chlorine gas treatment is widely and effectively used for municipal water purification/disinfection processes, making water safe for drinking and general public use.

Preparation: Ozone is prepared in the laboratory by passing a silent (non-sparking) electric discharge through dry oxygen gas, in a device called an ozoniser: $3O_2 \xrightarrow{\text{silent electric discharge}} 2O_3$

A silent discharge (rather than a sparking discharge) is specifically required, since the higher-energy spark discharge would tend to further decompose any ozone formed back into ordinary oxygen.

Chemical test: Ozone turns starch-iodide paper (or solution) blue-black — ozone oxidises iodide ions ($I^-$) to iodine ($I_2$), and the released iodine then forms a characteristic blue-black complex with the starch present, providing a simple, reliable confirmatory test for ozone's presence: $O_3+2KI+H_2O \rightarrow I_2+2KOH+O_2$

White phosphorus exists as discrete tetrahedral $P_4$ molecules, with each phosphorus atom bonded to three others at a bond angle of only 60° — significantly smaller than the ideal, unstrained bond angle expected for phosphorus's bonding (closer to 90-109°).

Red phosphorus has a polymeric, chain-like structure formed by linking together many $P_4$ tetrahedral units, with the strained 60° angles largely relieved through this polymeric linkage.

This significant angular strain present specifically in white phosphorus's small 60° bond angles makes its P-P bonds considerably weaker and more reactive (higher in stored potential/strain energy) compared to red phosphorus's more relaxed, unstrained polymeric structure — this strain-driven reactivity difference is precisely why white phosphorus is significantly more reactive (and notably more toxic/pyrophoric, spontaneously igniting in air) than the comparatively stable, less strained red phosphorus.

Thermal stability of Group 2 carbonates increases down the group: $BeCO_3 < MgCO_3 < CaCO_3 < SrCO_3 < BaCO_3$

This trend is explained by decreasing polarising power of the metal cation down the group (as ionic size increases and charge density decreases). Smaller, more highly charged cations (like $Be^{2+}$) more strongly polarise/distort the large carbonate anion's electron cloud, weakening the internal C-O bonding within carbonate and making it easier for the carbonate to decompose (releasing $CO_2$) upon heating.

Down the group, larger cations (like $Ba^{2+}$) have progressively weaker polarising power, causing less distortion of the carbonate ion, resulting in a more stable carbonate structure that requires a higher temperature to decompose — this decreasing cation polarising power down the group is precisely why thermal stability of the carbonates increases correspondingly down Group 2.

The first member of each p-block group (e.g., B, C, N, O, F in period 2) typically shows anomalous behaviour compared to the rest of its group, due to several combined factors:

1. Exceptionally small atomic/ionic size: The first member is significantly smaller than subsequent group members, resulting in unusually high ionisation enthalpy, electronegativity, and charge density, causing chemical behaviour that doesn't smoothly extrapolate from the rest of the group.

2. Absence of d-orbitals: Being in period 2, the first member has only s and p orbitals available in its valence shell, lacking accessible d-orbitals — this limits its maximum covalency (unable to expand its octet, unlike heavier group members which can utilise d-orbitals for expanded coordination, e.g., forming $PCl_5$ but not analogous $NCl_5$).

These combined factors (unusually small size and the specific unavailability of d-orbitals) are the primary general reasons why the first member of each p-block group so consistently shows distinctly anomalous chemical behaviour relative to its heavier group congeners.