p-Block Elements (Groups 15, 16, 17 and 18)

Group 15 elements (N, P, As, Sb, Bi) possess the valence shell electronic configuration ns2 np3, featuring a half-filled p-subshell that confers extra thermodynamic stability. Electronegativity, ionization enthalpy, and non-metallic character decrease down the group, whereas atomic radius, metallic character, and density increase progressively. The common oxidation states are -3, +3, and +5. However, due to the inert pair effect (the reluctance of the valence s-electrons to participate in bond formation owing to poor shielding by intervening d and f electrons), the stability of the +5 oxidation state decreases down the group while the stability of the +3 state increases: Bi(V) acts as an exceptionally powerful oxidizing agent (e.g., BiF₅).

Ammonia is manufactured industrially via Haber's Process by the direct reversible combination of dry dinitrogen and dihydrogen:

According to Le Chatelier's Principle, maximum equilibrium yield of NH₃ is favored by: (i) High pressure (~200 atm / 20 MPa) to shift equilibrium toward fewer gaseous moles; (ii) Optimum moderate temperature (~700 K / 450-500°C) balancing exothermic thermodynamics with kinetic activation; (iii) Iron oxide catalyst with promoter additives (K₂O and Al₂O₃).

Molecular Structure: The central nitrogen atom undergoes sp³ hybridization, possessing three bond pairs and one lone pair. The geometry is trigonal pyramidal with a bond angle of 107.8° (reduced from the ideal tetrahedral 109.5° due to lone pair-bond pair repulsion). The localized lone pair renders NH₃ a potent Lewis base, forming deep blue complex ions with Cu²⁺:

Manufactured industrially by Ostwald's Process in three sequential steps:

• Step 1 (Catalytic Oxidation): 4NH₃(g) + 5O₂(g) xrightarrow[500 K, 9 bar]{Pt/Rh gauge} 4NO(g) + 6H₂O(g)   (Rapid, ~95% conversion).
• Step 2 (NO Oxidation): 2NO(g) + O₂(g) ⇌ 2NO₂(g).
• Step 3 (Absorption): 3NO₂(g) + H₂O(l) → 2HNO₃(aq) + NO(g)   (NO recycled).

Oxidizing Action on Copper: Reaction products depend strictly on acid concentration:

Metal Passivity: Concentrated HNO₃ renders metals like Al, Cr, Fe passive due to the immediate formation of an impervious, non-reactive nanoscopic surface oxide film (e.g., Al₂O₃, Cr₂O₃).

Brown Ring Test Mechanism: In nitrate solution acidified with dilute H₂SO₄, freshly prepared FeSO₄ reduces NO₃^- to NO. The NO coordinates to pentaaquairon to yield the brown coordination complex:

In this complex, electron transfer occurs from NO to Fe(II), stabilizing the Fe⁺¹ oxidation state (3d⁷, high-spin, 3 unpaired electrons, magnetic moment μ ≈ 3.87 BM) with coordinated NO⁺ (nitrosonium ion).

Phosphorus displays multiple allotropic modifications with stark contrasts in stability and chemical reactivity:

Phosphine is prepared in the laboratory by heating white phosphorus with concentrated NaOH solution in an inert atmosphere of CO₂:

Spontaneous Inflammability: Pure PH₃ is non-inflammable. However, laboratory-prepared phosphine contains traces of diphosphine (P₂H₄) or P₄ vapor which spontaneously ignite upon contact with atmospheric oxygen, forming distinctive vortex rings of white P₄O₁₀ smoke:

Phosphorus Trichloride (PCl₃): Formed by passing dry chlorine over heated white phosphorus: P₄ + 6Cl₂ → 4PCl₃. Hydrolyzes violently with water to form phosphorous acid: PCl₃ + 3H₂O → H₃PO₃ + 3HCl. Geometry is pyramidal (sp³, 3 bp + 1 lp).

Phosphorus Pentachloride (PCl₅): Prepared by reaction with excess chlorine: P₄ + 10Cl₂ → 4PCl₅. In the vapor/gaseous phase, PCl₅ exhibits trigonal bipyramidal geometry (sp³d hybridization).

Group 16 elements (O, S, Se, Te, Po) are termed chalcogens (ore formers), having the valence electronic configuration ns2 np4. Ionization enthalpy is lower than Group 15 counterparts due to extra electron repulsion in the doubly occupied np_x orbital. Oxygen exhibits anomalous behavior: it forms pπ-pπ double bonds (O=O), exists as a paramagnetic diatomic gas (explained by Molecular Orbital Theory with 2 unpaired electrons in degenerate π*2p orbitals), has a maximum covalency of 4 (typically 2), and lacks vacant d-orbitals (hence OF₆ cannot exist, whereas SF₆ is an inert, octahedral gas).

Ozone is an allotrope of oxygen formed by passing a slow, dry stream of oxygen through a silent electric discharge:

A silent electric discharge prevents local overheating which would decompose ozone back to oxygen.

Thermodynamic Instability: Decomposition of ozone is spontaneous and exothermic: 2O₃(g) → 3O₂(g), where ΔH is negative and ΔS is positive (increase in moles). Thus, ΔG = ΔH - TΔS is heavily negative, making high concentrations of ozone potentially explosive.

Potent Oxidizing Action: Ozone releases nascent oxygen readily: O₃ → O₂ + [O]:

• Oxidizes black lead sulfide to white lead sulfate: PbS + 4O₃ → PbSO₄ + 4O₂.
• Oxidizes acidified potassium iodide to iodine: 2KI + H₂O + O₃ → 2KOH + I₂ + O₂.
• Tailing of mercury: Ozone oxidizes mercury to suboxide: 2Hg + O₃ → Hg₂O + O₂, causing mercury to lose its meniscus and stick to glass walls.

Structure of Ozone: Ozone has an angular/bent geometry (O-O-O bond angle = 117°) with two resonance hybrid structures where both O-O bond lengths are identical (128 pm), intermediate between single (148 pm) and double (121 pm) bond lengths.

Elemental sulfur forms puckered S₈ crown rings with an S-S-S bond angle of 107°:

• Rhombic Sulfur (α-sulfur): Yellow, stable below 368.6 K (95.6°C), density 2.06 g/cm³, prepared by evaporating roll sulfur solution in CS₂.
• Monoclinic Sulfur (β-sulfur): Needle-shaped, stable above 368.6 K (95.6°C), density 1.98 g/cm³, prepared by melting rhombic sulfur and cooling.
• Transition Temperature (368.6 K / 95.6°C): The exact temperature at which both rhombic and monoclinic allotropes coexist in dynamic equilibrium.
• Paramagnetism of S₂ in Vapor Phase: Above 1000 K, sulfur vapor exists predominantly as diatomic S₂ molecules, which are paramagnetic with 2 unpaired electrons in antibonding π* orbitals, behaving exactly analogous to O₂.

SO₂ is a pungent, choking gas prepared by roasting sulfide ores or burning sulfur: S + O₂ → SO₂. In aqueous solution, it forms sulfurous acid: SO₂ + H₂O ⇌ H₂SO₃.

Reducing Action: In the presence of moisture, SO₂ acts as a powerful reducing agent:

Sulfuric acid (H₂SO₄) is manufactured on a massive industrial scale via the Contact Process in four distinct technological stages:

• Stage 1 (Production of SO₂): Burning sulfur or roasting iron pyrites: S + O₂ → SO₂ or 4FeS₂ + 11O₂ → 2Fe₂O₃ + 8SO₂. The SO₂ gas is rigorously purified to remove dust, moisture, and arsenic impurities (As₂O₃ poisons the catalyst).
• Stage 2 (Catalytic Oxidation of SO₂ to SO₃ — Crucial Step):
  2SO₂(g) + O₂(g) ⇌ 2SO₃(g)   [ΔH° = -196.6 kJ/mol   (Exothermic)]
  Optimum Conditions: (i) Temperature: 720-770 K (450-500°C) — moderate temperature because the reaction is exothermic; (ii) Pressure: 1-2 bar (slight positive pressure to aid gas flow; high pressure favors SO₃ but moderate pressure gives ~98% yield economically); (iii) Catalyst: Vanadium pentoxide (V₂O₅), highly resistant to poisoning.
• Stage 3 (Absorption of SO₃ in conc. H₂SO₄): SO₃ is absorbed in 98% H₂SO₄ to produce Oleum (Pyrosulfuric acid): SO₃ + H₂SO₄ → H₂S₂O₇.
  Note: SO₃ is never dissolved directly in water because it forms an unmanageable dense acid mist/fog of H₂SO₄ that does not easily condense.
• Stage 4 (Dilution of Oleum): Oleum is diluted with calculated quantities of water to yield desired concentrations (e.g., 98% pure commercial grade): H₂S₂O₇ + H₂O → 2H₂SO₄.

Sulfuric acid owes its diverse industrial versatility to four characteristic properties:

• 1. Low Volatility: Displaces more volatile acids from their salts: 2NaCl + H₂SO₄ → Na₂SO₄ + 2HCl↑.
• 2. Strong Dibasic Acid Behavior: Ionizes in two steps:
  H₂SO₄(aq) + H₂O(l) → H₃O⁺(aq) + HSO₄⁻(aq)   [K_{a1} > 10   (Very Large)]
HSO₄⁻(aq) + H₂O(l) ⇌ H₃O⁺(aq) + SO₄²⁻(aq)   [K_{a2} = 1.2 × 10⁻²]
  Because K_a1 ≫ K_a2, H₂SO₄ behaves as a strong acid primarily in its first dissociation.
• 3. Powerful Dehydrating Agent: Removes elements of water with intense affinity:
  Charring of Cane Sugar: C₁₂H₂₂O₁₁ xrightarrow{conc. H₂SO₄} 12C(s) + 11H₂O (leaves black spongy carbon).
  Dehydration of Formic Acid: HCOOH xrightarrow{conc. H₂SO₄} CO↑ + H₂O.
• 4. Moderately Strong Oxidizing Agent: Concentrated hot H₂SO₄ oxidizes non-metals and metals:
  C + 2H₂SO₄(conc) → CO₂ + 2SO₂ + 2H₂O
S + 2H₂SO₄(conc) → 3SO₂ + 2H₂O
Cu + 2H₂SO₄(conc) xrightarrow{Δ} CuSO₄ + SO₂↑ + 2H₂O

Group 17 elements (F, Cl, Br, I, At) are known as halogens (salt formers), with valence electronic configuration ns2 np5, requiring just one electron to complete an octet. Fluorine shows distinctive anomalies: it exhibits only the -1 oxidation state (highest electronegativity 4.0, lacks vacant d-orbitals), whereas Cl, Br, and I exhibit oxidation states of -1, +1, +3, +5, and +7.

Halogens form four series of oxoacids: Hypohalous (HOX), Halous (HXO₂), Halic (HXO₃), and Perhalic (HXO₄). Fluorine forms only one oxoacid, HOF (fluoric(I) acid or hypofluorous acid).

Acidic Strength Order of Chlorine Oxoacids:

Explanation: (1) As the oxidation state of the central halogen increases, electron density is pulled away from the O-H bond, weakening it and facilitating proton release. (2) In the conjugate base series (ClO₄^-, ClO₃^-, ClO₂^-, ClO^-), the negative charge in ClO₄⁻ is delocalized over four equivalent electronegative oxygen atoms via resonance, providing maximal stabilization.

Halogens react among themselves to form binary covalent compounds called interhalogens of types XX', XX'₃, XX'₅, and XX'₇ (where X is larger and more electropositive than X'):

Why are Interhalogens More Reactive Than Parent Halogens? In general, interhalogens (XX') are chemically more reactive than dihalogens (X₂) except fluorine (F₂). This is because the heterogeneous X-X' covalent bond is polarized and significantly weaker than the homogeneous non-polar X-X bond in pure halogens.

Group 18 elements (He, Ne, Ar, Kr, Xe, Rn) possess completely filled valence shells (ns2 np6, except He with 1s2). Due to this stable octet configuration, they exhibit exceptionally high first ionization enthalpies and large positive electron gain enthalpies, remaining chemically unreactive under normal conditions.

Xenon reacts directly with fluorine in nickel vessels under controlled stoichiometry, temperature, and pressure:

• Xe + F₂ xrightarrow[673 K, 1 bar]{Xe in excess (2:1)} XeF₂ (Xenon difluoride, colorless solid).
• Xe + 2F₂ xrightarrow[873 K, 7 bar]{1:5 molar ratio} XeF₄ (Xenon tetrafluoride).
• Xe + 3F₂ xrightarrow[573 K, 60-70 bar]{1:20 molar ratio} XeF₆ (Xenon hexafluoride).

Hydrolysis Reactions (Complete vs Partial):

• Helium (He): (1) Used as a diluent for oxygen in deep-sea diving tanks (heliox) because unlike nitrogen, helium has exceptionally low solubility in blood and lipids, preventing the decompression illness known as 'the bends'. (2) Used to fill meteorological balloons due to low density and non-inflammability (unlike H₂). (3) Liquid helium (b.p. 4.2 K) serves as an essential cryogenic refrigerant for superconducting magnets in Nuclear Magnetic Resonance (NMR) and Medical Magnetic Resonance Imaging (MRI).
• Neon (Ne): Filled in discharge tubes and neon glow signs for vibrant orange-red advertising displays; neon lamps produce light that easily penetrates fog.
• Argon (Ar): Provides an inert gas atmosphere in high-temperature metallurgical processes such as arc welding of aluminum and stainless steel, and filling incandescent light bulbs to inhibit tungsten filament evaporation.
• Krypton & Xenon (Kr, Xe): Utilized in high-speed photographic flash lamps, airport runway approach beacons, and specialized laser systems.
• Radon (Rn): Radioactive noble gas employed in radiotherapy for malignant tumor treatment and geological earthquake tracking.