Comprehensive Study Notes on p-Block Chemistry (Class 11 & 12)

Fundamental Periodic Properties and Trends in p-Block Elements

• Classification of p-Block Elements:

  • General outer electronic configuration spans from $ns^2 np^1$ to $ns^2 np^6$ (with Helium possessing a unique $1s^2$ configuration).

  • The block contains a mix of metals, non-metals, metalloids, and synthetically produced radioactive elements.

  • Group-wise Valence Shell Configurations:

  • Group 13 (Boron family, icosagens): $ns^2 np^1$

  • Group 14 (Carbon family, crystallogens): $ns^2 np^2$

  • Group 15 (Nitrogen family, pnictogens): $ns^2 np^3$

  • Group 16 (Oxygen family, chalcogens): $ns^2 np^4$

  • Group 17 (Halogens): $ns^2 np^5$

  • Group 18 (Noble gases, aerogens): $ns^2 np^6$

  • Core Electron Filling and Inner Shells:

  • Heavier members of Period 4 contain completely filled $3d^{10}$ subshells.

  • Members of Period 5 contain filled $4d^{10}$ subshells.

  • Members of Period 6 contain filled $4f^{14} 5d^{10}$ subshells, giving rise to relativistic and shielding phenomena.

  • Liquid-to-Solid Expansion Anomaly:

  • Gallium ($Ga$), Germanium ($Ge$), and Bismuth ($Bi$) exhibit the rare physical property where their liquid phase expands upon solidification, meaning the density of the solid state is lower than that of the molten liquid phase (analogous to water and ice).

• Atomic Radii Trends and Contraction Anomalies:

  • Atomic radii typically increase down a group due to the addition of principal energy levels ($n$).

  • Group 13 Radius Inversion:

  • The observed order of atomic radius is:       $B < Ga < Al < In < Tl$

  • Atomic radius of Gallium ($135\,pm$) is smaller than that of Aluminium ($143\,pm$).

  • Origin: Transition mental/d-block contraction. Gallium is preceded by ten $3d$ transition elements. The intervening $3d^{10}$ electrons have poor shielding ability, causing a substantial increase in effective nuclear charge ($Z_{\text{eff}}$), which pulls the outer $4s$ and $4p$ electrons inward.

• Ionisation Energy (IE) Trends:

  • Generally decreases down each group due to increasing atomic size and screening effects.

  • Group 13 Ionisation Energy Order:

  • Experimental order: $B > Tl > Ga > Al > In$

  • Across Group 13, the IE curve displays a characteristic "W"-shaped profile.

  • Boron ($B$) possesses the highest ionisation energy in the group, whereas Indium ($In$) has the lowest.

  • $Tl > Ga > Al > In$ arises because:

    • Gallium has higher IE than Aluminium due to the $3d^{10}$ transition contraction.

    • Thallium has a higher IE than Gallium, Aluminium, and Indium due to lanthanoid contraction ($4f^{14}$ subshell provides negligible shielding, resulting in a large $Z_{\text{eff}}$).

  • Extremes in the Periodic Table:

  • Highest ionisation energy across the entire periodic table: Helium ($He$).

  • Lowest ionisation energy in the periodic table: Caesium ($Cs$).

  • Metallic Character Correlation:

  • As ionisation energy decreases down a group, the tendency to lose valence electrons increases, directly elevating metallic (electropositive) character.

• Electron Gain Enthalpy ($\Delta H_{eg}$) and Electron Affinity:

  • Endothermic Electron Gain Enthalpies ($\Delta H_{eg} > 0$):

  • Group 2 elements ($Be$, $Mg$) due to completely filled, stable $ns^2$ subshells.

  • Nitrogen ($N$) due to a stable half-filled $2p^3$ subshell ($N + e^- + \text{energy} \rightarrow N^-$).

  • Group 18 Noble gases due to fully saturated octets/duplets ($ns^2 np^6$).

  • Any mono-negative anion adding a second electron ($X^- + e^- \rightarrow X^{2-}$) due to immense electrostatic electron-electron repulsion.

  • Exothermic Electron Gain Enthalpies ($\Delta H_{eg} < 0$):

  • Characteristic of elements with high electron affinity, releasing energy upon electron capture.

  • Halogen order of negative magnitude: $|Cl| > |F| > |Br| > |I|$

  • Chlorine ($Cl$) has the most negative $\Delta H_{eg}$ across the periodic table ($-349\,kJ\,mol^{-1}$).

  • Fluorine ($F$) possesses a lower magnitude ($-328\,kJ\,mol^{-1}$) because the added electron enters a compact $2p$ subshell where it encounters strong inter-electronic repulsions.

  • Noble Gas Positivity Sequence:

  • The element with the most positive $\Delta H_{eg}$ in the entire periodic table is Neon ($Ne$).

  • Relative order of positive $\Delta H_{eg}$ among noble gases: $Ne > Ar = Kr > Xe > Rn > He$

• Electronegativity Trends:

  • Group 13 Pauling Electronegativity Values:

  • Values: $B\,(2.0)$, $Al\,(1.5)$, $Ga\,(1.6)$, $In\,(1.7)$, $Tl\,(1.8)$.

  • Order: $B > Tl > In > Ga > Al$.

  • Electronegativity decreases initially from $B$ to $Al$, but steadily increases from $Ga$ to $Tl$ due to ineffective screening by intervening $d$ and $f$ orbitals.

  • Group 14 Pauling Electronegativity Values:

  • Values: $C\,(2.5)$, $Si\,(1.8)$, $Ge\,(1.8)$, $Sn\,(1.8)$, $Pb\,(1.9)$.

  • Order: $C > Pb > Si \approx Ge \approx Sn$.

Inert Pair Effect and Variable Oxidation States

• Definition and Mechanical Basis:

  • The inert pair effect is the reluctance of the valence $s$-electrons ($6s^2$ in Period 6) to participate in chemical bonding.

  • Causal Mechanism: Caused by poor screening of nuclear charge by the intervening $4f^{14}$ (lanthanoid contraction) and $5d^{10}$ electrons. This ineffective shielding increases $Z_{\text{eff}}$, tightly binding the outermost $6s^2$ electron pair to the nucleus.

  • Consequence: The lower oxidation state ($\text{Group Oxidation State} - 2$) becomes increasingly stable down the group, while the higher group oxidation state decreases in stability.

• Group-Wise Stability Comparisons:

  • Group 13 (+1 vs +3):

  • Stability of +3 state: $B^{3+} > Al^{3+} > Ga^{3+} > In^{3+} > Tl^{3+}$

  • Stability of +1 state: $B^+ < Al^+ < Ga^+ < In^+ < Tl^+$

  • $Tl^+$ is significantly more stable than $Tl^{3+}$.

  • $Tl^{3+}$ behaves as a potent oxidising agent: $Tl^{3+} + 2e^- \rightarrow Tl^+$

  • Group 14 (+2 vs +4):

  • Stability of +4 state: $C^{4+} > Si^{4+} > Ge^{4+} > Sn^{4+} > Pb^{4+}$

  • Stability of +2 state: $C^{2+} < Si^{2+} < Ge^{2+} < Sn^{2+} < Pb^{2+}$

  • In Lead, the +2 state is thermodynamically favored ($Pb^{2+} > Pb^{4+}$).

  • In Tin, the +4 state is thermodynamically favored ($Sn^{4+} > Sn^{2+}$).

  • Redox Activity:

    • $Pb^{4+}$ compounds act as strong oxidising agents ($Pb^{4+} + 2e^- \rightarrow Pb^{2+}$).

    • $Sn^{2+}$ compounds act as strong reducing agents ($Sn^{2+} \rightarrow Sn^{4+} + 2e^-$).

    • Thermal decomposition of $Pb(IV)$ species yields $Pb(II)$ species:         $PbCl_4 \xrightarrow{\Delta} PbCl_2 + Cl_2\uparrow$         $PbO_2 \xrightarrow{\Delta} PbO + \frac{1}{2} O_2\uparrow$

    • Tin(II) species readily undergo oxidation:         $SnCl_2 + Cl_2 \rightarrow SnCl_4$         $SnO + \frac{1}{2} O_2 \rightarrow SnO_2$

  • Group 15 (+3 vs +5):

  • Stability of +5 state: $N^{5+} > P^{5+} > As^{5+} > Sb^{5+} > Bi^{5+}$

  • Stability of +3 state: $N^{3+} < P^{3+} < As^{3+} < Sb^{3+} < Bi^{3+}$

  • $Bi^{5+}$ is an exceptionally strong oxidising agent. The only well-characterised stable $Bi(V)$ compound is $BiF_5$, formed because of the small size, maximum electronegativity, and superior oxidising power of Fluorine ($F_2$).

• Existence vs. Non-Existence of Polyhalides and Lattice Structures:

  • Non-existence of $PbI_4$:

  • $PbI_4$ does not exist under standard conditions because $Pb^{4+}$ is a strong oxidising agent and $I^-$ is a strong reducing agent. A redox reaction takes place instantly:       $Pb^{4+} + 2I^- \rightarrow Pb^{2+} + I_2$

  • Hence, $PbI_2$ exists stably, but $PbI_4$ cannot form.

  • Structural Nature of $TlI_3$:

  • $TlI_3$ is isomorphous to Caesium triiodide ($CsI_3$).

  • It does not contain Thallium in the +3 oxidation state ($Tl^{3+} + 3I^-$ does not exist due to the reduction of $Tl^{3+}$ by $I^-$).

  • It is formulated as an ionic lattice containing Thallium(I) and the linear triiodide anion: $Tl^+ [I_3]^-$

  • Mixed Valence Gallium Dichloride ($GaCl_2$):

  • Gallium dichloride is diamagnetic and does not contain $Ga(II)$.

  • It exists as a dimer $Ga_2Cl_4$, structurally composed of a mixed-valence coordination complex: $Ga^+ [GaCl_4]^-$ ($GaCl \cdot GaCl_3$), containing Gallium in +1 and +3 oxidation states.

Anomalous Behavior of Second-Period Elements

• Distinctive Characteristics:

  • Unusually small atomic and ionic radii.

  • High electronegativity and high ionisation enthalpy.

  • Absolute absence of low-lying, vacant $d$-orbitals in their valence shell ($n = 2$ contains only $2s$ and $2p$ orbitals).

• Covalency Restrictions:

  • Second-period elements have a maximum covalency of 4 (utilising one $2s$ and three $2p$ orbitals).

  • Boron can form $[BF_4]^-$ (covalency = 4, oxidation state = +3), but cannot expand its octet to form $[BF_6]^{3-}$.

  • Aluminium, possessing vacant $3d$-orbitals, can expand its coordination shell to form $[AlF_6]^{3-}$.

  • Nitrogen can achieve a maximum covalency of 4 ($NH_4^+$) and cannot form pentahalides like $NCl_5$ or $NF_5$. In contrast, Phosphorus forms $PCl_5$, $PF_5$, and $[PF_6]^-$.

• Steric Constraints in Heavier Congeners:

  • Although Silicon has vacant $3d$-orbitals, $[SiCl_6]^{2-}$ does not exist.

  • Mechanisms:

  • Six large chloride ions ($Cl^-$) cannot be sterically accommodated around the relatively small $Si^{4+}$ central ion.

  • Strong electrostatic lone pair-lone pair repulsions occur between the adjacent bulky chloride ligands.

  • In contrast, the smaller fluoride ion forms stable $[SiF_6]^{2-}$.

• Multi-Bonding Propensity ($p\pi-p\pi$ vs. $d\pi-p\pi$):

  • Second-period elements uniquely form strong, stable $p\pi-p\pi$ multiple bonds with themselves and other small, highly electronegative atoms:     $C=C,\quad C\equiv C,\quad C=O,\quad C\equiv N,\quad N\equiv N,\quad O=O$

  • Heavier p-block elements have diffuse, expanded $p$-orbitals that cannot effectively overlap sideways. Consequently, multiple bonding involving heavier elements relies predominantly on $d\pi-p\pi$ or $d\pi-d\pi$ interactions.

  • Example: In trimethylamine oxide ($(CH_3)_3N \rightarrow O$), the $N-O$ linkage is a dative single bond because Nitrogen lacks $d$-orbitals to form a double bond. In contrast, trimethylphosphine oxide ($(CH_3)_3P=O$) contains a short, strong $P=O$ double bond involving $p\pi(O) \rightarrow d\pi(P)$ back-donation.

Bond Dissociation Energy and Catenation Properties

• Single-Bond Dissociation Energies (BDE):

  • Group 14 Single-Bond Strengths:     $C-C\,(348\,kJ\,mol^{-1}) > Si-Si\,(297\,kJ\,mol^{-1}) > Ge-Ge\,(260\,kJ\,mol^{-1}) > Sn-Sn\,(240\,kJ\,mol^{-1})$

  • Group 15 Single-Bond Strengths:

  • $P-P > N-N$

  • The single $N-N$ bond is substantially weaker than the $P-P$ single bond due to strong inter-electronic repulsion between the non-bonding lone pairs localized on the small, tightly packed Nitrogen atoms.

  • Conversely, the triple bond in dinitrogen ($N\equiv N$) is exceptionally strong ($941.4\,kJ\,mol^{-1}$), explaining the inertness of molecular nitrogen at ambient temperatures.

  • Group 16 Single-Bond Strengths:

  • $S-S > O-O$

  • Due to lone-pair repulsion on the compact Oxygen atoms in the peroxide linkage ($-O-O-$), the bond energy is low relative to the catenated $-S-S-$ linkage.

  • Group 17 Diatomic Bond Dissociation Energies:

  • Observed sequence: $Cl_2 > Br_2 > F_2 > I_2$

  • Fluorine ($F_2$) exhibits an anomalously low dissociation energy compared to Chlorine and Bromine due to severe lone-pair repulsions across its short ($142\,pm$) internuclear distance.

• Catenation:

  • Definition: The innate tendency of an element to form continuous covalent chains of identical atoms.

  • Catenation depends primarily on the magnitude of the single covalent bond enthalpy ($E-E$).

  • Group 14 Trend:     $C \gg Si > Ge \approx Sn$

  • Carbon exhibits the highest catenation propensity among all elements, forming rings, chains, and lattices.

  • Lead ($Pb$) does not display catenation.

  • Group 15 Trend:

  • Phosphorus has a significantly greater catenation tendency than Nitrogen, forming elemental chains and rings ($P_4$, $(P_4)_n$).

  • Bismuth does not catenate and exists purely through metallic bonding in its elemental state.

Allotropes of Carbon, Phosphorus, and Sulphur

• Allotropes of Carbon:

  • Diamond:

  • Crystalline Network: Giant 3-D rigid covalent network where every Carbon atom is $sp^3$ hybridized and tetrahedrally bonded to four others.

  • Bond Metrics: All $C-C$ bond lengths are identical at $154\,pm$ ($1.54\,\text{Å}$).

  • Properties: Hardest naturally occurring material, non-conductor of electricity (zero mobile electrons), excellent thermal conductor (efficient phonon transmission through rigid lattice), high melting point ($> 4000\,K$).

  • Graphite:

  • Layered Structure: Composed of 2-D planar hexagonal sheets of $sp^2$ hybridized Carbon atoms.

  • Bond Metrics: In-plane $C-C$ bond length is $141.5\,pm$ ($1.415\,\text{Å}$), shorter than diamond due to delocalized $\pi$-resonance. The interlayer separation is $340\,pm$.

  • Bonding and Lubrication: Layers are held by weak van der Waals forces, enabling them to shear over one another smoothly. Used as a dry lubricant at elevated operating temperatures where liquid grease or oil oxidizes.

  • Conductivity: The fourth valence electron occupies a delocalized $\pi$-system covering the layer. Conducts electricity parallel to the graphene layers, but acts as a semiconductor perpendicular to them.

  • Thermodynamic Reference: Thermodynamically, graphite is the most stable allotrope of Carbon at standard temperature and pressure; its standard enthalpy of formation is zero:       $\Delta H_f^\circ(\text{graphite}) = 0\,kJ\,mol^{-1}$       $\Delta H_f^\circ(\text{diamond}) = +1.9\,kJ\,mol^{-1}$

  • Fullerenes ($C_{60}$ - Buckminsterfullerene):

  • Synthesis: Prepared by evaporating graphite using an electric arc in an inert atmosphere of Helium or Argon.

  • Topology: Geodesic soccer-ball cage consisting of 60 vertices, 90 edges ($C-C$ bonds), and 32 faces.

  • Ring Count: Exactly 20 six-membered rings (hexagons) and 12 five-membered rings (pentagons).

  • Fusion Rules:

    • A six-membered ring is fused to both six-membered rings and five-membered rings.

    • A five-membered ring is fused exclusively to six-membered rings (pentagon-pentagon fusion is structurally prohibited).

  • Hybridization: Each Carbon is $sp^2$ hybridized with three $\sigma$-bonds and one delocalized $\pi$-electron. Two distinct bond lengths exist: single bonds between pentagon-hexagon junctions ($\sim 143.5\,pm$) and double bonds between hexagon-hexagon junctions ($\sim 138.3\,pm$).

  • It represents the only pure allotrope of carbon, having a smooth cage with no uncoordinated "dangling" bonds.

• Allotropes of Phosphorus:

  • White Phosphorus ($P_4$):

  • Structure: Discrete tetrahedral $P_4$ molecules.

  • Bond Metrics and Angle Strain: Each P atom is $sp^3$ hybridized with one lone pair. The theoretical tetrahedral angle ($109.5^\circ$) is compressed to an internal bond angle of $60^\circ$. This produces immense angular strain ($\approx 49.5^\circ$), leading to weak, bent ("banana") $P-P$ bonds.

  • Molecular Features: Contains 6 $P-P$ single bonds and 4 non-bonding lone pairs.

  • Chemical Behavior: Translucent, white, waxy solid. Highly poisonous. Insoluble in water, highly soluble in carbon disulfide ($CS_2$). Shows chemiluminescence (glows greenish in air). Ignites spontaneously in air at $\sim 303\,K$:       $P_4 + 5O_2 \rightarrow P_4O_{10} \quad (\text{dense white fumes})$

  • Disproportionation with Base: Dissolves in boiling aqueous Sodium Hydroxide under an inert atmosphere ($CO_2$) to produce phosphine gas:       $P_4 + 3NaOH + 3H_2O \rightarrow PH_3\uparrow + 3NaH_2PO_2$

  • Red Phosphorus ($(P_4)_n$):

  • Structure: Polymeric chain of $P_4$ tetrahedra linked by breaking one $P-P$ bond per tetrahedron.

  • Preparation: Formed by heating white phosphorus at $573\,K$ in an inert atmosphere for several days.

  • Properties: Iron-grey lustre, non-poisonous, odorless, insoluble in both water and $CS_2$. Much less reactive than white phosphorus; does not chemiluminesce.

  • Black Phosphorus:

  • Exists in two main crystalline modifications: $\alpha$-black and $\beta$-black.

  • $\alpha$-Black Phosphorus: Formed by heating red phosphorus in a sealed tube at $803\,K$. Forms opaque monoclinic or rhombohedral crystals; does not oxidise in air.

  • $\beta$-Black Phosphorus: Formed by heating white phosphorus at $473\,K$ under high hydrostatic pressure ($\sim 12000\,\text{atm}$). Has a corrugated, puckered sheet structure and conducts electricity like graphite.

  • Thermodynamic Stability Order:       $\text{Black Phosphorus} > \text{Red Phosphorus} > \text{White Phosphorus}$

• Allotropes of Sulphur:

  • Rhombic Sulphur ($\alpha$-Sulphur):

  • Stable crystalline form below $369\,K$.

  • Yellow color, composed of puckered, crown-shaped $S_8$ rings.

  • Insoluble in water, dissolves readily in $CS_2$.

  • Monoclinic Sulphur ($\beta$-Sulphur):

  • Stable crystalline form above $369\,K$.

  • Needle-shaped crystals, also containing puckered $S_8$ rings.

  • Transition Temperature ($369\,K$ or $96^\circ C$):

  • Both $\alpha$ and $\beta$ allotropes are in dynamic thermodynamic equilibrium at this temperature:       $\alpha\text{-Sulphur} \xrightleftharpoons{369\,K} \beta\text{-Sulphur}$

  • $S_6$ Ring:

  • Adopts a non-planar "chair" conformation.

  • High-Temperature Vapour Phase ($S_2$ Form):

  • At temperatures of approximately $1000\,K$, sulphur predominantly vaporises as diatomic $S_2$.

  • Paramagnetism: According to Molecular Orbital Theory (MOT), $S_2$, like $O_2$, has two unpaired electrons in its doubly degenerate antibonding $\pi^*$ molecular orbitals ($\pi^*_{2p_x}$, $\pi^*_{2p_y}$; or for sulphur, $\pi^*_{3p_x}$, $\pi^*_{3p_y}$), making it paramagnetic. The $S_8$ ring is diamagnetic.

Physical Property Trends and Group Anomalies

• Boiling Point (BP) Systematics:

  • Groups 13 and 14: Decrease monotonically down the group:

  • Group 13: $B > Al > Ga > In > Tl$

  • Group 14: $C > Si > Ge > Sn > Pb$

  • Groups 15, 16, 17, and 18: Increase down the group (driven by increasing molecular mass and polarizability, which strengthen London dispersion forces):

  • Group 18: $He < Ne < Ar < Kr < Xe < Rn$

• Melting Point (MP) Systematics:

  • Group 13: $B > Al > Tl > In > Ga$

  • Gallium has an unusually low melting point ($303\,K$ or $30^\circ C$) because it crystallizes into discrete $Ga_2$ diatomic units rather than a continuous metal lattice. It remains liquid over a very broad temperature range, making it ideal for high-temperature thermometers.

  • Group 14: $C > Si > Ge > Pb > Sn$

  • Group 15: $N < P < As > Sb > Bi$

  • Groups 16, 17, and 18: Monotonically increase down the group.

• Density Trends:

  • Density ($d = m/V$) generally increases down any group in the p-block.

  • Anomaly: Carbon (as diamond, $d \approx 3.51\,g\,cm^{-3}$) has a higher density than Silicon ($d \approx 2.33\,g\,cm^{-3}$) due to its tightly packed, short $C-C$ covalent lattice.

• Boron Isotopes and Nuclear Applications:

  • Naturally occurs as two stable isotopes:

  • $^{10}B$ (natural abundance $\approx 19\%$)

  • $^{11}B$ (natural abundance $\approx 81\%$)

  • The $^{10}B$ isotope has an extremely high capture cross-section for thermal neutrons.

  • Metal borides and boron fibers are widely used as neutron-absorbing protective shields and control rods in nuclear reactors, and as light structural composites for military aircraft and bullet-proof body armor.

Chemistry and Structural Systematics of Oxoacids

• Nomenclature Nomenclature Framework:

  • "-ic" Acid: Central non-metal in its standard high/group oxidation state ($H_3PO_4$, $H_2SO_4$, $HNO_3$, $HClO_3$).

  • "-ous" Acid: Derived by removing one oxygen from the "-ic" acid ($[-O]$):

  • $H_3PO_4 - [O] = H_3PO_3$ (Phosphorous acid)

  • $H_2SO_4 - [O] = H_2SO_3$ (Sulphurous acid)

  • $HNO_3 - [O] = HNO_2$ (Nitrous acid)

  • $HClO_3 - [O] = HClO_2$ (Chlorous acid)

  • "hypo-…-ous" Acid: Derived by removing one oxygen from the "-ous" acid ($[-O]$):

  • $H_3PO_3 - [O] = H_3PO_2$ (Hypophosphorous acid)

  • $HNO_2 - [O] = \frac{1}{2} H_2N_2O_2$ (Hyponitrous acid, $H_2N_2O_2$)

  • $HClO_2 - [O] = HClO$ (Hypochlorous acid)

  • "per-…-ic" or "peroxy-…-ic" Acid: Derived by adding one oxygen to the "-ic" acid ($[+O]$):

  • $H_2SO_4 + [O] = H_2SO_5$ (Peroxomonosulphuric acid or Caro's acid)

  • $HNO_3 + [O] = HNO_4$ (Peroxonitric acid)

  • $HClO_3 + [O] = HClO_4$ (Perchloric acid)

  • "pyro-…-ic" Acid: Dimerisation of two "-ic" acid molecules with the elimination of one water molecule ($2\times\text{"-ic acid"} - H_2O$):

  • $2H_3PO_4 - H_2O = H_4P_2O_7$ (Pyrophosphoric acid / Diphosphoric acid)

  • $2H_2SO_4 - H_2O = H_2S_2O_7$ (Pyrosulphuric acid / Oleum)

  • "pyro-…-ous" Acid: Dimerisation of two "-ous" acid molecules minus one water molecule ($2\times\text{"-ous acid"} - H_2O$):

  • $2H_2SO_3 - H_2O = H_2S_2O_5$ (Pyrosulphurous acid)

  • $2H_3PO_3 - H_2O = H_4P_2O_5$ (Pyrophosphorous acid)

  • "meta-…-ic" Acid: Formed by the direct dehydration of a single molecule of an "-ic" acid ($\text{"-ic acid"} - H_2O$):

  • $H_3PO_4 - H_2O = HPO_3$ (Metaphosphoric acid)

  • $H_3BO_3 - H_2O = HBO_2$ (Metaboric acid)

• Acids Containing a Single Central Atom:

  • Phosphorus Oxoacids:

  • Every Phosphorus oxoacid has a tetrahedral coordination geometry with at least one $P=O$ bond and at least one $P-OH$ bond.

  • Any hydrogen directly attached to phosphorus ($P-H$ bond) is non-ionisable and acts as a strong reducing agent.

  • Only hydrogens in $P-OH$ groups are ionisable (determining acid basicity/proticity).

  • Orthophosphoric Acid ($H_3PO_4$):

    • Oxidation state: +5

    • Composition: One $P=O$, three $P-OH$

    • Basicity: 3 (Tribasic)

    • Reducing property: Non-reducing

  • Phosphorous Acid ($H_3PO_3$):

    • Oxidation state: +3

    • Composition: One $P=O$, two $P-OH$, one $P-H$

    • Basicity: 2 (Dibasic)

    • Reducing property: Moderate reducing agent

  • Hypophosphorous Acid ($H_3PO_2$):

    • Oxidation state: +1

    • Composition: One $P=O$, one $P-OH$, two $P-H$

    • Basicity: 1 (Monobasic)

    • Reducing property: Strong reducing agent

  • Sulphur Oxoacids with a Single S Atom:

  • Sulphuric Acid ($H_2SO_4$): S is in +6 OS; structure has two $S=O$ double bonds and two $S-OH$ single bonds (dibasic).

  • Sulphurous Acid ($H_2SO_3$): S is in +4 OS; structure has one $S=O$ double bond, two $S-OH$ single bonds, and one lone pair on S (dibasic).

  • Caro's Acid ($H_2SO_5$ - Peroxomonosulphuric acid): S is in +6 OS; contains two $S=O$ bonds, one $S-OH$ bond, and one peroxomonosulphate link ($S-O-O-H$).

  • Silicic Acid ($H_4SiO_4$ - Orthosilicic acid): Silicon is $sp^3$ hybridized, bonded to four $-OH$ groups via single covalent bonds.

• Structural Rules for Acids with Two Central Atoms ($A_2$):

  • Calculate the Theoretical Oxidation State (TOS) assuming standard valencies ($H = +1$, $O = -2$):     $\text{TOS} = \frac{2(\text{number of } O) - (\text{number of } H)}{2}$

  • Rule 1: If $\text{TOS} < \text{Max OS}$:

  • The two central atoms are linked directly via an $A-A$ covalent bond (no bridging oxygen).

  • Hypophosphoric Acid ($H_4P_2O_6$): TOS = +4, Max OS of P = +5. Contains a direct $P-P$ bond. Structure: $(HO)_2(O)P-P(O)(OH)_2$.

  • Dithionic Acid ($H_2S_2O_6$): TOS = +5, Max OS of S = +6. Contains a direct $S-S$ bond. Structure: $HO(O)_2S-S(O)_2OH$.

  • Dithionous Acid ($H_2S_2O_4$): TOS = +3, Max OS of S = +6. Contains an $S-S$ bond: $HO(O)S-S(O)OH$.

  • Oxalic Acid ($H_2C_2O_4$): TOS = +3, Max OS of C = +4. Contains a direct $C-C$ bond: $HO(O)C-C(O)OH$.

  • Rule 2: If $\text{TOS} = \text{Max OS}$:

  • The central atoms are bridged by a single oxygen atom, forming an $A-O-A$ ether-like linkage.

  • Pyrophosphoric Acid ($H_4P_2O_7$): TOS = +5 = Max OS of P. Contains a $P-O-P$ bridge. Structure: $(HO)_2(O)P-O-P(O)(OH)_2$.

    • Contains 2 $P=O$ bonds, 4 $P-OH$ bonds, 1 $P-O-P$ bridge.

    • Total $\sigma$-bonds = 12; total $\pi$-bonds = 2. Ratio of $\sigma/\pi = 12/2 = 6$.

  • Pyrosulphuric Acid / Oleum ($H_2S_2O_7$): TOS = +6 = Max OS of S. Contains an $S-O-S$ bridge. Structure: $HO(O)_2S-O-S(O)_2OH$.

  • Pyrosilicic Acid ($H_6Si_2O_7$): TOS = +4 = Max OS of Si. Contains a $Si-O-Si$ bridge: $(HO)_3Si-O-Si(OH)_3$.

  • Rule 3: If $\text{TOS} > \text{Max OS}$:

  • The central atoms are linked by a peroxide bridge ($A-O-O-A$), and both central atoms remain in their maximum possible oxidation state.

  • Marshall's Acid ($H_2S_2O_8$ - Peroxodisulphuric acid): TOS = +7, but Max OS of S = +6. Contains an $S-O-O-S$ bridge. Actual OS on both Sulphur atoms is +6.

  • Peroxydiphosphoric Acid ($H_4P_2O_8$): TOS = +6, but Max OS of P = +5. Contains a $P-O-O-P$ bridge.

• Polythionic Acids ($H_2S_xO_6$, where $x = 3\text{ to }5$):

  • General formulation: $HO-SO_2-(S)_{x-2}-SO_2-OH$

  • Oxidation States of Sulphur:

  • The two terminal Sulphur atoms are bonded to oxygens and have an oxidation state of +5 each.

  • The intermediate, catenated Sulphur atoms ($x - 2$ atoms) are bonded exclusively to other sulphur atoms and carry an oxidation state of 0.

  • Thus, the oxidation states present are 0 and +5.

Classification, Structure, and Acidity of p-Block Oxides

• Acid-Base Character Trends:

  • As the oxidation state of an element increases, its oxide becomes increasingly acidic.

  • Down a group, basic character increases and acidic character decreases as metallic radius increases and electronegativity drops.

  • Group 13 Oxides:

  • $B_2O_3$: Strictly acidic (dissolves in water to form weak orthoboric acid, $H_3BO_3$).

  • $Al_2O_3$ and $Ga_2O_3$: Amphoteric (react with both acids and strong bases).

  • $In_2O_3$: Weakly basic.

  • $Tl_2O_3$ and $Tl_2O$: Strongly basic.

  • Group 14 Oxides:

  • $CO$: Neutral.

  • $CO_2$, $SiO_2$, $GeO_2$: Acidic.

  • $SnO_2$, $PbO_2$, $SnO$, $PbO$: Amphoteric.

  • Mixed Oxides: Red Lead ($Pb_3O_4 = 2PbO \cdot PbO_2$) and $Pb_2O_3 = PbO \cdot PbO_2$.

  • Group 15 Oxides:

  • Nitrogen and Phosphorus oxides in higher oxidation states (+5, +4, +3) are acidic.

  • Arsenic and Antimony trioxides ($As_2O_3$, $Sb_2O_3$) are amphoteric.

  • Bismuth trioxide ($Bi_2O_3$) is predominantly basic.

• Oxides of Nitrogen:

  • Dinitrogen Monoxide ($N_2O$, Nitrous oxide):

  • Oxidation state: +1

  • Chemical nature: Neutral gas

  • Structure: Linear, diamagnetic, isoelectronic with $CO_2$. Resonance forms involve $:N\equiv N\rightarrow O:$ and $N=N=O$.

  • Nitrogen Monoxide ($NO$, Nitric oxide):

  • Oxidation state: +2

  • Chemical nature: Neutral gas

  • Structure: Monomeric gas is paramagnetic due to an odd electron in its $\pi^*$ orbital. In liquid and solid states, it dimerises to diamagnetic $N_2O_2$.

  • Dinitrogen Trioxide ($N_2O_3$):

  • Oxidation state: +3

  • Chemical nature: Acidic, intensely blue solid at low temperatures.

  • Structure: Planar molecule with an unsymmetrical $N-N$ bond ($O=N-NO_2$).

  • Nitrogen Dioxide ($NO_2$):

  • Oxidation state: +4

  • Chemical nature: Acidic, brown, toxic gas.

  • Structure: Paramagnetic odd-electron molecule with a single unpaired electron on Nitrogen. Readily dimerises via an exothermic equilibrium to colourless, diamagnetic Dinitrogen tetroxide:       $2NO_2\text{ (brown, paramagnetic)} \xrightleftharpoons{\text{cool}} N_2O_4\text{ (colourless, diamagnetic)}$

  • $N_2O_4$ contains a long, weak planar $N-N$ covalent bond ($178\,pm$).

  • Dinitrogen Pentoxide ($N_2O_5$):

  • Oxidation state: +5

  • Chemical nature: Strongly acidic, anhydride of nitric acid ($2HNO_3 - H_2O = N_2O_5$).

  • Structure in gas/liquid state: Planar covalent molecule with an oxygen bridge: $O_2N-O-NO_2$. Does not contain an $N-N$ bond.

  • Structure in crystalline solid state: Exists as an ionic solid composed of linear nitronium cations and planar nitrate anions: $[NO_2]^+ [NO_3]^-$.

• Phosphorus Oxides ($P_4O_6$ and $P_4O_{10}$):

  • Derived by inserting Oxygen atoms into the edges and vertices of a $P_4$ tetrahedron.

  • Tetraphosphorus Hexaoxide ($P_4O_6$): Six bridging Oxygen atoms are inserted into each of the 6 edges of the $P_4$ tetrahedron ($P-O-P$ linkages). Contains 12 $P-O$ single bonds and 4 non-bonding lone pairs on the Phosphorus atoms.

  • Tetraphosphorus Decaoxide ($P_4O_{10}$): In addition to the 6 bridging oxygens along the edges, 4 terminal oxygen atoms are coordinatively double-bonded to each of the 4 Phosphorus vertices ($P=O$). Contains 12 $P-O$ bridging single bonds and 4 $P=O$ double bonds (consisting of $p\pi-d\pi$ overlap).

Hydrides of Groups 15, 16, and 17: Physical and Chemical Trends

• Physical States and Hydrogen Bonding:

  • Water ($H_2O$) is a liquid at room temperature ($\text{b.p.} = 373\,K$), whereas Hydrogen Sulphide ($H_2S$) is a gas ($\text{b.p.} = 213\,K$). Oxygen's small size and high electronegativity facilitate extensive 3D intermolecular hydrogen bonding (4 hydrogen bonds per $H_2O$ molecule in ice).

  • Hydrogen Fluoride ($HF$) is a low-boiling liquid ($\text{b.p.} = 293\,K$) due to intermolecular hydrogen bonding forming zig-zag chains, whereas $HCl$, $HBr$, and $HI$ are gases at room temperature.

• Boiling Point Sequences:

  • Group 14 Hydrides: Monotonically increases with molecular weight:     $CH_4 < SiH_4 < GeH_4 < SnH_4$

  • Group 15 Hydrides:     $PH_3 < AsH_3 < NH_3 < SbH_3 < BiH_3$

  • Ammonia ($NH_3$) has an elevated boiling point due to intermolecular hydrogen bonding, placing it above $PH_3$ and $AsH_3$, but below $SbH_3$ and $BiH_3$, where strong London dispersion forces dominate.

  • Group 16 Hydrides:     $H_2S < H_2Se < H_2Te < H_2O$

  • $H_2O$ has the highest boiling point in the group due to its dense hydrogen-bonding network.

  • Group 17 Hydrides:     $HCl < HBr < HI < HF$

  • $HF$ possesses the highest boiling point among the halogen acids due to the strong $F-H\cdots F$ hydrogen bond.

  • Relative Boiling Point of Key Hydrogen-Bonded Hydrides:     $H_2O > HF > NH_3$

  • Although the individual $H\cdots F$ bond is stronger than the $H\cdots O$ bond, each water molecule can form an average of four hydrogen bonds, creating a dense 3D network, whereas $HF$ forms linear chains with only two hydrogen bonds per molecule.

• Periodic Trends in Binary Hydrides Down Groups 15, 16, and 17:

  • Bond Length ($E-H$): Increases down the group as the covalent radius of the central atom expands:

  • Group 15: $N-H < P-H < As-H < Sb-H < Bi-H$

  • Group 16: $O-H < S-H < Se-H < Te-H$

  • Group 17: $F-H < Cl-H < Br-H < I-H$

  • Bond Dissociation Energy (BDE) and Thermal Stability: Decreases down the group as orbital overlap becomes more diffuse:

  • Group 15: $NH_3 > PH_3 > AsH_3 > SbH_3 > BiH_3$

  • Group 16: $H_2O > H_2S > H_2Se > H_2Te$

  • Group 17: $HF > HCl > HBr > HI$

  • Acidic Strength: Increases down the group. As the $E-H$ bond dissociation energy drops, proton ($H^+$) release in polar solutions becomes thermodynamically favored:

  • Group 16: $H_2O < H_2S < H_2Se < H_2Te$

  • Group 17: $HF < HCl < HBr < HI$

  • Reducing Character: Increases down the group as the stability of the hydride decreases:

  • Group 15: $NH_3 < PH_3 < AsH_3 < SbH_3 < BiH_3$ ($BiH_3$ is the strongest reducing agent among the pnictogen hydrides).

  • Group 16: $H_2O < H_2S < H_2Se < H_2Te$

  • Basic Strength: Decreases down Group 15:     $NH_3 > PH_3 > AsH_3 > SbH_3 > BiH_3$

  • As the central atom size increases, the non-bonding lone pair occupies a more diffuse orbital with lower electron density, drastically reducing its availability for protonation.

Back Bonding: Principles, Overlap Mechanics, and Molecular Manifestations

• Mechanics and Governing Conditions:

  • Back bonding is a type of resonance involving coordinate $\pi$-bonding, formed by the sideways overlap of a filled atomic orbital on one atom with an empty, symmetry-compatible orbital on an adjacent atom.

  • Necessary Conditions:

  1. One participating atom must belong to the second period (possessing compact, localized $2p$ orbitals).

  2. The partner atom must belong to either the second or third period (with low-lying vacant $2p$, $3p$, or $3d$ orbitals).

  • Relative $\pi$-Overlap Strength:     $2p\pi-2p\pi > 2p\pi-3d\pi \approx 2p\pi-3p\pi \gg 3d\pi-3d\pi$

  • Overlap strength decreases as the sum of principal quantum numbers ($n_1 + n_2$) increases.

  • Consequences on Molecular Geometry:

  • Increase in bond order, causing bond length contraction and higher bond dissociation energy.

  • Alteration of central atom hybridization and conversion of pyramidal geometry to planar.

  • Attenuation or neutralization of classic Lewis acidity or basicity.

• Lewis Acidity of Boron Trihalides ($BX_3$):

  • Boron trihalides are electron-deficient monomers with 6 valence electrons on Boron, leaving an unhybridized, vacant $2p_z$ orbital perpendicular to the planar $sp^2$ framework.

  • Halogens donate a non-bonding lone pair from their filled $p$-orbital into the empty $2p_z$ orbital of Boron ($p\pi-p\pi$ back-donation).

  • Overlap Efficiency:

  • In $BF_3$: Boron ($2p$) and Fluorine ($2p$) form an efficient $2p\pi-2p\pi$ bond due to comparable orbital sizes.

  • In $BCl_3$: Boron ($2p$) and Chlorine ($3p$) have a less effective $2p\pi-3p\pi$ overlap.

  • In $BBr_3$ ($2p\pi-4p\pi$) and $BI_3$ ($2p\pi-5p\pi$), overlap efficiency drops further.

  • Observed Lewis Acidity Order:     $BF_3 < BCl_3 < BBr_3 < BI_3$

  • In $BF_3$, back bonding satisfies Boron's electron deficiency to the greatest extent, making it the weakest Lewis acid. $BI_3$ receives minimal back-donation and acts as the strongest Lewis acid.

  • Bond Contraction in $BF_3$ vs. Adducts:

  • The $B-F$ bond length in monomeric $BF_3$ is $1.30\,\text{Å}$, significantly shorter than the single covalent bond length ($1.52\,\text{Å}$) due to partial double-bond character from back bonding.

  • In the trimethylamine adduct ($(CH_3)_3N \cdot BF_3$) or tetrafluoroborate ($[BF_4]^-$), the empty $2p$ orbital on Boron is occupied by a coordinate $\sigma$-bond. Boron's hybridization converts from $sp^2$ to $sp^3$, eliminating back bonding. As a result, the $B-F$ bond length expands to $1.35\,\text{Å}$.

• Structure and Basicity of Silylamines vs. Alkylamines:

  • Trimethylamine ($N(CH_3)_3$):

  • Nitrogen is $sp^3$ hybridized.

  • Geometry: Trigonal pyramidal with an active lone pair in an $sp^3$ hybrid orbital.

  • Chemical nature: Acts as a strong classical Lewis base.

  • Trisilylamine ($N(SiH_3)_3$):

  • Silicon belongs to the third period and has vacant $3d$-orbitals.

  • The non-bonding lone pair on Nitrogen delocalizes into the vacant $3d$-orbitals of the three surrounding Silicon atoms through $2p\pi-3d\pi$ back bonding.

  • Steric Number ($SN$) of Nitrogen: $\text{Lone pairs} + \sigma\text{-bonds} = 0 + 3 = 3$.

  • Hybridization: Nitrogen converts to $sp^2$ planar geometry ($\angle Si-N-Si = 120^\circ$).

  • Chemical nature: Weak Lewis base because the lone pair is not localized on Nitrogen.

• Molecular Geometry of Isocyanates:

  • Methyl Isocyanate ($CH_3NCO$):

  • Carbon lacks accessible $d$-orbitals, so no back bonding occurs.

  • Nitrogen retains a localized lone pair: $SN = 1\text{ lone pair} + 2\,\sigma\text{-bonds} = 3$.

  • Hybridization: $sp^2$, resulting in a bent $C-N=C=O$ backbone ($\angle C-N=C \approx 125^\circ$).

  • Silyl Isocyanate ($SiH_3NCO$):

  • The lone pair on Nitrogen back-donates into the empty $3d$-orbitals of Silicon via a $2p\pi-3d\pi$ interaction.

  • Steric Number of Nitrogen: $SN = 0\text{ localized lone pairs} + 2\,\sigma\text{-bonds} = 2$.

  • Hybridization: $sp$, producing a strictly linear $Si-N=C=O$ geometry ($\angle Si-N=C = 180^\circ$).