Class 12 Chemistry Notes Chapter 7 (The p-block elements) – Chemistry-I Book

Alright students, let's focus on one of the most significant chapters for your exams from the Class 12 syllabus – Chapter 7, 'The p-Block Elements'. This block is unique because it contains metals, non-metals, and metalloids, showcasing a wide range of chemical behaviour. We will cover Groups 15 to 18 as per your syllabus. Pay close attention to the trends, structures, important reactions, and properties of key compounds.

The p-Block Elements (Groups 15-18)

General Introduction:

• Elements in which the last electron enters any of the three p-orbitals of their outermost shell are called p-block elements.
• Groups 13 to 18 constitute the p-block.
• General electronic configuration: ns²np¹⁻⁶ (except for He, which is 1s²).
• Properties are greatly influenced by atomic size, ionization enthalpy, electron gain enthalpy, and electronegativity.
• Presence of non-metals and metalloids is a key feature. Metallic character increases down the group, while non-metallic character increases across a period.

Group 15 Elements (Nitrogen Family)

• Elements: Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb), Bismuth (Bi).
• Electronic Configuration: ns²np³ (stable half-filled p-orbitals).
• Occurrence: N₂ (78% of air), Minerals (NaNO₃ - Chile saltpetre, KNO₃ - Indian saltpetre), Proteins. P (Apatite family, bones, DNA/RNA). As, Sb, Bi (Sulphide minerals).
• Trends:
  • Atomic/Ionic Radii: Increase down the group. Significant increase from N to P, smaller increase thereafter due to poor shielding by d (As, Sb) and f (Bi) orbitals.
  • Ionization Enthalpy (IE): Decreases down the group. Very high IE due to stable half-filled configuration. IE(Group 15) > IE(Group 14) & IE(Group 16) in the same period.
  • Electronegativity: Decreases down the group.
  • Physical Properties: N₂ is a gas, others are solids. Metallic character increases (N, P - non-metals; As, Sb - metalloids; Bi - metal). Boiling points increase down the group, melting point increases up to As and then decreases up to Bi. All show allotropy except Nitrogen and Bismuth (P-white, red, black; As-yellow, grey; Sb-yellow, grey).
  • Catenation: Tendency to link with itself. P >> N > As > Sb. N=N triple bond is strong, but N-N single bond is weak due to repulsion between lone pairs on small N atoms.
• Chemical Properties:
  • Oxidation States: Common are -3, +3, +5. Tendency to show -3 state decreases down the group. Stability of +5 state decreases, and +3 state increases down the group due to the Inert Pair Effect (reluctance of ns² electrons to participate in bonding). Bi predominantly shows +3. Nitrogen exhibits oxidation states from -3 to +5.
  • Anomalous Behaviour of Nitrogen: Due to its small size, high electronegativity, high ionization enthalpy, and absence of d-orbitals. N can form pπ-pπ multiple bonds (N≡N), has a maximum covalency of 4 (cannot form NF₅, NCl₅), while others can expand their covalency (e.g., PCl₅, PCl₆⁻).
  • Reactivity towards Hydrogen (Hydrides - EH₃): All form EH₃ (NH₃, PH₃, AsH₃, SbH₃, BiH₃). Structure: Pyramidal. Basic character decreases (NH₃ > PH₃ > AsH₃ > SbH₃ > BiH₃). Thermal stability decreases down the group. Reducing character increases down the group. Bond angle decreases down the group (lone pair-bond pair repulsion effect). NH₃ has H-bonding.
  • Reactivity towards Oxygen (Oxides): Form E₂O₃ and E₂O₅ types. Oxides in higher oxidation state are more acidic. Acidic character decreases down the group (N₂O₅ > P₂O₅ > As₂O₅ > Sb₂O₅ > Bi₂O₅). Oxides of N, P are acidic; As, Sb are amphoteric; Bi is basic. Nitrogen forms numerous oxides (N₂O, NO, N₂O₃, NO₂, N₂O₄, N₂O₅).
  • Reactivity towards Halogens (Halides): Form EX₃ (pyramidal) and EX₅ (trigonal bipyramidal) types. EX₅ are more covalent than EX₃. All trihalides except N are stable. NX₃ (except NF₃) are unstable. Pentahalides are less stable than trihalides; stability decreases down the group (PF₅ > AsF₅ > SbF₅; BiF₅ is unknown). PCl₅ exists as [PCl₄]⁺[PCl₆]⁻ in solid state. Trihalides (except BiF₃) are covalent and hydrolyse in water (PCl₃ + 3H₂O → H₃PO₃ + 3HCl). Pentahalides also hydrolyse (PCl₅ + H₂O → POCl₃ + 2HCl; POCl₃ + 3H₂O → H₃PO₄ + 3HCl). N cannot form pentahalides due to absence of d-orbitals.
  • Reactivity towards Metals: Form binary compounds showing -3 oxidation state (e.g., Ca₃N₂, Ca₃P₂, Na₃As).
• Important Compounds:
  • Dinitrogen (N₂): Lab Prep: NH₄Cl(aq) + NaNO₂(aq) → N₂(g) + 2H₂O(l) + NaCl(aq). Commercial: Liquefaction and fractional distillation of air. Properties: Colourless, odourless, tasteless, non-toxic gas. Very inert at room temperature due to high N≡N bond enthalpy. Reacts with metals at high temp (Li₃N, Mg₃N₂). Reacts with H₂ at ~773K (Haber's process). Uses: Manufacturing NH₃, inert atmosphere, cryosurgery, refrigerant.
  • Ammonia (NH₃): Prep: Haber's Process (N₂ + 3H₂ ⇌ 2NH₃; High P ~200 atm, Temp ~700K, Catalyst: FeO with K₂O, Al₂O₃). Lab Prep: Heating ammonium salts with base (e.g., NH₄Cl + Ca(OH)₂ → CaCl₂ + 2NH₃ + 2H₂O). Structure: Pyramidal, N is sp³ hybridized. Properties: Colourless gas, pungent odour. Highly soluble in water (forms NH₄OH), basic (Lewis base). Forms complexes (e.g., [Cu(NH₃)₄]²⁺, [Ag(NH₃)₂]⁺). Uses: Fertilizers (urea, ammonium nitrate), manufacture of HNO₃, refrigerant.
  • Nitric Acid (HNO₃): Prep: Ostwald's Process (Catalytic oxidation of NH₃: 4NH₃ + 5O₂ → 4NO + 6H₂O [Pt/Rh catalyst, 500K, 9 bar]; 2NO + O₂ → 2NO₂; 3NO₂ + H₂O → 2HNO₃ + NO). Lab Prep: NaNO₃ + H₂SO₄(conc) → NaHSO₄ + HNO₃. Properties: Colourless liquid. Strong oxidizing agent (reacts with most metals except noble metals like Au, Pt). Action depends on concentration, temperature, and nature of metal. Cu(dil HNO₃) → Cu(NO₃)₂ + NO; Cu(conc HNO₃) → Cu(NO₃)₂ + NO₂. Non-metals oxidized (I₂ → HIO₃, C → CO₂, S → H₂SO₄, P₄ → H₃PO₄). Brown Ring Test for nitrates. Structure: Planar molecule. Uses: Fertilizers, explosives (TNT), dyes, perfumes, pickling of stainless steel, oxidizer in rocket fuels.
  • Phosphorus (P): Allotropes: White P (discrete P₄ tetrahedral units, reactive, poisonous, chemiluminescent, soluble in CS₂), Red P (polymeric chain structure, less reactive, non-poisonous, insoluble in CS₂), Black P (α and β forms, most stable, layered structure, semiconductor).
  • Phosphine (PH₃): Prep: Ca₃P₂ + 6H₂O → 3Ca(OH)₂ + 2PH₃. Lab Prep: Heating white P with conc. NaOH (P₄ + 3NaOH + 3H₂O → PH₃ + 3NaH₂PO₂). Properties: Colourless gas, rotten fish smell, highly poisonous. Weakly basic (Lewis base). Uses: Holme's signals, smoke screens.
  • Phosphorus Halides: PCl₃ (Prep: P₄ + 6Cl₂ → 4PCl₃) and PCl₅ (Prep: P₄ + 10Cl₂ → 4PCl₅). Both used as chlorinating agents in organic chemistry. PCl₃ (Pyramidal), PCl₅ (Trigonal bipyramidal in gas/liquid; [PCl₄]⁺[PCl₆]⁻ in solid). Hydrolysis reactions are important.
  • Oxoacids of Phosphorus: H₃PO₂ (Hypophosphorous acid, monobasic), H₃PO₃ (Orthophosphorous acid, dibasic), H₃PO₄ (Orthophosphoric acid, tribasic), H₄P₂O₇ (Pyrophosphoric acid, tetrabasic). Structures involve tetrahedral P atom(s) with at least one P=O bond and one P-OH bond. Acids with P-H bonds act as reducing agents (e.g., H₃PO₂, H₃PO₃). Basicity depends on the number of P-OH groups.

Group 16 Elements (Oxygen Family - Chalcogens)

• Elements: Oxygen (O), Sulfur (S), Selenium (Se), Tellurium (Te), Polonium (Po - radioactive).
• Electronic Configuration: ns²np⁴.
• Occurrence: O (Most abundant element on Earth, 21% air by volume). S (Sulphates - gypsum CaSO₄.2H₂O, epsom MgSO₄.7H₂O; Sulphides - galena PbS, zinc blende ZnS). Se, Te (Tellurides in sulphide ores). Po (Decay product of Th, U).
• Trends:
  • Atomic/Ionic Radii: Increase down the group.
  • Ionization Enthalpy (IE): Decreases down the group. IE(Group 16) < IE(Group 15) due to stable half-filled p-orbitals in Group 15.
  • Electron Gain Enthalpy: O has less negative EGE than S due to small size and high electron density. Trend: S > Se > Te > Po > O.
  • Electronegativity: Decreases down the group. O is the second most electronegative element.
  • Physical Properties: O₂ is gas, others solids. Metallic character increases (O, S - non-metals; Se, Te - metalloids; Po - metal). Melting/Boiling points increase down the group. All show allotropy (O-O₂, O₃; S-Rhombic, Monoclinic).
  • Catenation: S >> Se > Te > O. S-S bond is strong.
• Chemical Properties:
  • Oxidation States: Common are -2, +2, +4, +6. Oxygen mostly shows -2 (except OF₂ (+2), peroxides (-1), superoxides (-1/2)). Stability of -2 state decreases down the group. Stability of +6 decreases, +4 increases down the group (Inert Pair Effect). O cannot show > +2 due to absence of d-orbitals. S, Se, Te show +4, +6 due to available d-orbitals.
  • Anomalous Behaviour of Oxygen: Due to small size, high electronegativity, absence of d-orbitals. Forms strong pπ-pπ bonds (O=O). Exists as diatomic molecule (O₂), others form puckered rings (S₈) or chains. H-bonding in H₂O leads to high boiling point. Maximum covalency is 2 (can extend to 4 in H₃O⁺), while others can show up to 6 (e.g., SF₆).
  • Reactivity towards Hydrogen (Hydrides - H₂E): Form H₂E (H₂O, H₂S, H₂Se, H₂Te). Structure: Bent/Angular. Acidic character increases (H₂O < H₂S < H₂Se < H₂Te). Thermal stability decreases (H₂O > H₂S > H₂Se > H₂Te). Reducing character increases down the group. Bond angle decreases down the group (except H₂O). H₂O is liquid (H-bonding), others are gases.
  • Reactivity towards Oxygen (Oxides): Form EO₂ (e.g., SO₂) and EO₃ (e.g., SO₃). O₃ is ozone. SO₂ is gas, SeO₂ is solid. Reducing property of dioxides decreases (SO₂ > SeO₂ > TeO₂). SO₂, SeO₂, TeO₂ are acidic. SO₃, SeO₃, TeO₃ are also acidic. Acidity decreases down the group.
  • Reactivity towards Halogens (Halides): Form EX₂, EX₄, EX₆. Stability: F⁻ > Cl⁻ > Br⁻ > I⁻. Hexafluorides (SF₆, SeF₆, TeF₆) are stable, gaseous, octahedral. SF₆ is exceptionally stable (steric hindrance, kinetically inert). Tetrafluorides (SF₄, SeF₄, TeF₄) are gases/liquids/solids, see-saw shape (sp³d). Dihalides (except O) are sp³ hybridized, bent shape. Monohalides are dimeric (e.g., S₂F₂, S₂Cl₂).
• Important Compounds:
  • Dioxygen (O₂): Lab Prep: Heating KClO₃ (MnO₂ catalyst), thermal decomposition of oxides (Ag₂O, HgO), H₂O₂ decomposition. Commercial: Liquefaction and fractional distillation of air, electrolysis of water. Properties: Colourless, odourless gas. Paramagnetic (due to unpaired electrons in π* antibonding MOs). Reacts with most metals and non-metals (Oxidation). Uses: Respiration, combustion, oxy-acetylene welding, steel manufacture.
  • Ozone (O₃): Allotrope of Oxygen. Prep: Passing silent electric discharge through pure, cold, dry O₂ (3O₂ ⇌ 2O₃). Structure: Angular/Bent molecule, resonance hybrid. Properties: Pale blue gas, characteristic smell. Thermodynamically unstable compared to O₂. Powerful oxidizing agent (oxidizes PbS to PbSO₄, I⁻ to I₂). Used as germicide, disinfectant, sterilizing water, bleaching oils/flour. Depletion by CFCs in stratosphere is a concern.
  • Sulphur (S): Allotropes: Rhombic (α-Sulphur, stable at room temp, S₈ puckered rings), Monoclinic (β-Sulphur, stable above 369K, S₈ puckered rings). Both soluble in CS₂. S₆ ring form also exists. Paramagnetism of S₂ vapour at high temp (~1000K) similar to O₂.
  • Sulphur Dioxide (SO₂): Prep: Burning S in air (S + O₂ → SO₂), Roasting sulphide ores (FeS₂ + O₂ → Fe₂O₃ + SO₂). Lab Prep: Na₂SO₃ + dil H₂SO₄ → Na₂SO₄ + H₂O + SO₂. Structure: Angular/Bent. Properties: Colourless gas, pungent smell. Highly soluble in water (forms H₂SO₃). Acts as reducing agent (with KMnO₄, K₂Cr₂O₇) and oxidizing agent (with H₂S). Bleaching action (temporary, due to reduction). Uses: Refining petroleum/sugar, bleaching wool/silk, disinfectant, manufacture of H₂SO₄, NaHSO₃.
  • Sulphuric Acid (H₂SO₄): "King of Chemicals". Prep: Contact Process (Burning S or sulphide ores → SO₂; Catalytic oxidation of SO₂ to SO₃: 2SO₂ + O₂ ⇌ 2SO₃ [V₂O₅ catalyst, 720K, 2 bar]; Absorption of SO₃ in conc. H₂SO₄ to form Oleum (H₂S₂O₇); Dilution of Oleum with water → H₂SO₄). Properties: Colourless, dense, oily liquid. Low volatility, strong acidic character, strong affinity for water (powerful dehydrating agent - chars carbohydrates), strong oxidizing agent (oxidizes C, S, P, metals like Cu, Zn). Uses: Fertilizers (ammonium sulphate, superphosphate), petroleum refining, detergents, paints, pigments, storage batteries, manufacture of chemicals.
  • Oxoacids of Sulphur: H₂SO₃ (Sulphurous acid), H₂SO₄ (Sulphuric acid), H₂S₂O₇ (Pyrosulphuric acid/Oleum), H₂S₂O₈ (Peroxodisulphuric acid/Marshall's acid), H₂S₂O₃ (Thiosulphuric acid). Structures involve tetrahedral S atom(s).

Group 17 Elements (Halogen Family)

• Elements: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At - radioactive).
• Electronic Configuration: ns²np⁵.
• Occurrence: Highly reactive, found as halides. Fluorspar (CaF₂), Cryolite (Na₃AlF₆). Common salt (NaCl), Carnallite (KCl.MgCl₂.6H₂O). Sea water (bromides, iodides). Sea weeds (I).
• Trends:
  • Atomic/Ionic Radii: Smallest in their periods. Increase down the group.
  • Ionization Enthalpy (IE): Very high. Decreases down the group.
  • Electron Gain Enthalpy (EGE): Maximum negative EGE in the periodic table. Anomaly: EGE(Cl) > EGE(F) due to very small size and high inter-electronic repulsion in F. Order: Cl > F > Br > I.
  • Electronegativity: Highest in their periods. F is the most electronegative element. Decreases down the group (F > Cl > Br > I).
  • Physical Properties: Exist as diatomic molecules (F₂, Cl₂, Br₂, I₂). State at room temp: F₂, Cl₂ (gases); Br₂ (liquid); I₂ (solid). Coloured due to absorption of radiation in visible region (F₂-yellow, Cl₂-greenish yellow, Br₂-red, I₂-violet). Bond Dissociation Enthalpy Anomaly: Cl₂ > Br₂ > F₂ > I₂ (due to lone pair-lone pair repulsion in small F₂ molecule). Soluble in organic solvents, sparingly soluble in water (except F₂ which reacts).
• Chemical Properties:
  • Oxidation States: F always shows -1. Others show -1, +1, +3, +5, +7 (due to d-orbitals). Higher oxidation states are realized when halogens combine with more electronegative elements like F or O.
  • Reactivity: Highly reactive. Reactivity decreases down the group. Strong oxidizing agents. Oxidizing power: F₂ > Cl₂ > Br₂ > I₂. F₂ oxidizes water. Cl₂ and Br₂ react with water to form corresponding hydrohalic and hypohalous acids. Halogens displace less electronegative halogens from their salt solutions (e.g., Cl₂ + 2KBr → 2KCl + Br₂).
  • Anomalous Behaviour of Fluorine: Due to small size, highest electronegativity, low F-F bond dissociation enthalpy, absence of d-orbitals. Shows only -1 oxidation state. Forms strongest H-bonds (in HF). Highest standard electrode potential (strongest oxidizing agent). HF is a liquid (b.p. 293K), others are gases.
  • Reactivity towards Hydrogen (Hydrides - HX): Form HX (HF, HCl, HBr, HI). Covalent compounds. Acidic strength increases: HF < HCl < HBr < HI. Thermal stability decreases: HF > HCl > HBr > HI. Boiling point trend: HF >> HI > HBr > HCl (due to H-bonding in HF).
  • Reactivity towards Oxygen (Oxides): Oxides are generally unstable. Fluorine forms OF₂ (+2 state for O) and O₂F₂ (-1 state for O). Oxides of Cl (Cl₂O, ClO₂, Cl₂O₆, Cl₂O₇), Br (Br₂O, BrO₂, BrO₃), I (I₂O₄, I₂O₅, I₂O₇) are known. Higher oxides are more stable. Oxides of Cl, Br, I are powerful oxidizing agents and acidic. I₂O₅ is used in estimation of CO.
  • Reactivity towards Metals: React readily to form metal halides (e.g., NaCl, MgBr₂, AlCl₃). Ionic character decreases: MF > MCl > MBr > MI. Halides in higher oxidation states are more covalent (e.g., SnCl₄ > SnCl₂).
  • Reactivity towards other Halogens (Interhalogen Compounds): React amongst themselves to form interhalogen compounds (Types: XY, XY₃, XY₅, XY₇ where X is larger/less electronegative halogen, Y is smaller/more electronegative halogen). Examples: ClF, BrF₃, IF₅, IF₇. They are more reactive than constituent halogens (except F₂) because X-Y bond is weaker than X-X or Y-Y bond (except F-F). Structures based on VSEPR theory (ClF-linear, ClF₃/BrF₃-bent T-shape, IF₅-square pyramidal, IF₇-pentagonal bipyramidal). Undergo hydrolysis.
• Important Compounds:
  • Chlorine (Cl₂): Prep: Deacon's process (Oxidation of HCl by O₂ with CuCl₂ catalyst: 4HCl + O₂ → 2Cl₂ + 2H₂O). Lab Prep: MnO₂ + 4HCl(conc) → MnCl₂ + Cl₂ + 2H₂O; or heating KMnO₄/K₂Cr₂O₇ with conc HCl. Properties: Greenish-yellow gas, pungent suffocating odour. Soluble in water (forms HCl and HOCl). Strong oxidizing and bleaching agent (bleaching action due to HOCl → HCl + [O], permanent). Reacts with NH₃ (excess NH₃ gives N₂ + NH₄Cl; excess Cl₂ gives NCl₃). Reacts with NaOH (cold/dil → NaCl + NaOCl; hot/conc → NaCl + NaClO₃). Uses: Bleaching wood pulp, cotton, textiles; sterilizing drinking water; manufacture of dyes, drugs, organic compounds (CCl₄, DDT, refrigerants), poisonous gases (phosgene, tear gas).
  • Hydrogen Chloride (HCl): Lab Prep: NaCl + H₂SO₄(conc) → NaHSO₄ + HCl (at <420K); NaHSO₄ + NaCl → Na₂SO₄ + HCl (at >820K). Properties: Colourless pungent gas. Highly soluble in water (forms Hydrochloric acid). Strong acid. Reacts with metals, bases, salts. Forms aqua regia (conc HCl : conc HNO₃ = 3:1) which dissolves noble metals (Au, Pt). Uses: Manufacture of Cl₂, NH₄Cl, glucose from corn starch; extracting glue from bones; pickling steel; medicine; lab reagent.
  • Oxoacids of Halogens: Fluorine forms only HOF (Hypofluorous acid). Others form four series: Hypohalous acid (HOX), Halous acid (HXO₂ - only for Cl), Halic acid (HXO₃), Perhalic acid (HXO₄). Acid strength increases with oxidation state: HOX < HXO₂ < HXO₃ < HXO₄. Thermal stability also increases. Oxidizing power decreases: HOX > HXO₂ > HXO₃ > HXO₄. Structures involve central halogen atom.

Group 18 Elements (Noble Gases)

• Elements: Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), Radon (Rn - radioactive).
• Electronic Configuration: ns²np⁶ (except He: 1s²). Stable octet (duplet for He).
• Occurrence: All except Rn occur in atmosphere (Ar is major component ~1%). He found in natural gas. Rn is decay product of Radium.
• Trends:
  • Atomic Radii: Largest in their respective periods (van der Waals radii). Increase down the group.
  • Ionization Enthalpy (IE): Extremely high due to stable configuration. Decreases down the group.
  • Electron Gain Enthalpy: Large positive values (energy needed to add electron).
  • Physical Properties: Monoatomic gases. Colourless, odourless, tasteless. Sparingly soluble in water. Very low melting and boiling points (weak van der Waals forces). M.P./B.P. increase down the group. Can be liquefied at low temperatures.
• Chemical Properties:
  • Inertness: Generally unreactive due to stable configuration, high IE, positive EGE.
  • Discovery of Compounds: Neil Bartlett (1962) prepared the first compound O₂⁺[PtF₆]⁻. Noting that IE(O₂) ≈ IE(Xe), he reacted Xe with PtF₆ to form the first noble gas compound Xe⁺[PtF₆]⁻.
  • Xenon Compounds: Xe reacts directly with F₂ under different conditions to form XeF₂, XeF₄, XeF₆. Kr forms only KrF₂. No true compounds of He, Ne, Ar known.
    • Xenon Fluorides: XeF₂, XeF₄, XeF₆. Colourless crystalline solids. Powerful fluorinating agents. Hydrolysed by water (XeF₂ slowly, XeF₄ and XeF₆ violently).
      • XeF₂ + 2H₂O → 2Xe + 4HF + O₂ (slow)
      • 6XeF₄ + 12H₂O → 4Xe + 2XeO₃ + 24HF + 3O₂
      • XeF₆ + 3H₂O → XeO₃ + 6HF (complete hydrolysis)
      • XeF₆ + H₂O → XeOF₄ + 2HF (partial hydrolysis)
      • XeF₆ + 2H₂O → XeO₂F₂ + 4HF (partial hydrolysis)
    • Structures (VSEPR): XeF₂ (Linear), XeF₄ (Square planar), XeF₆ (Distorted octahedral).
    • Xenon Oxides: XeO₃ (Colourless explosive solid, pyramidal structure). XeO₄ (Unstable gas, tetrahedral structure).
    • Xenon Oxyfluorides: XeOF₄ (Colourless volatile liquid, square pyramidal structure), XeO₂F₂ (Colourless solid, see-saw structure).
• Uses:
  • Helium (He): Non-inflammable, light gas - filling balloons/airships. Cryogenic agent (low temp experiments, MRI magnets). Diluent for oxygen in deep sea diving (less soluble in blood than N₂).
  • Neon (Ne): Discharge tubes, fluorescent bulbs ('neon signs'). Botanical gardens, green houses.
  • Argon (Ar): Inert atmosphere in high-temp metallurgy, arc welding. Filling electric bulbs (prevents filament evaporation).
  • Krypton (Kr) & Xenon (Xe): Special purpose light bulbs, high-intensity lamps, lasers.
  • Radon (Rn): Radiotherapy for cancer treatment.

This covers the core concepts of p-Block elements (Groups 15-18). Remember to focus on trends, reasons for anomalous behaviour of the first element in each group, structures (using VSEPR), preparation methods, and key reactions/properties of important compounds like Ammonia, Nitric Acid, Ozone, Sulphuric Acid, Chlorine, HCl, and Xenon compounds.

Multiple Choice Questions (MCQs)

1. Which of the following hydrides has the lowest boiling point?
   (a) H₂O
   (b) H₂S
   (c) H₂Se
   (d) H₂Te

2. The stability of +5 oxidation state decreases down the group in Group 15 due to:
   (a) Increasing metallic character
   (b) Increasing ionization enthalpy
   (c) Inert pair effect
   (d) Decreasing electronegativity

3. The structure of XeF₄ molecule is:
   (a) Linear
   (b) Pyramidal
   (c) Square planar
   (d) Tetrahedral

4. Which of the following oxoacids of phosphorus is dibasic?
   (a) H₃PO₂ (Hypophosphorous acid)
   (b) H₃PO₃ (Orthophosphorous acid)
   (c) H₃PO₄ (Orthophosphoric acid)
   (d) H₄P₂O₇ (Pyrophosphoric acid)

5. The correct order of electron gain enthalpy (with negative sign) for halogens is:
   (a) F > Cl > Br > I
   (b) Cl > F > Br > I
   (c) Cl > Br > F > I
   (d) I > Br > Cl > F

6. Which gas is produced when copper reacts with concentrated nitric acid?
   (a) NO
   (b) N₂O
   (c) NO₂
   (d) N₂

7. SF₆ is exceptionally stable and inert primarily due to:
   (a) High S-F bond energy
   (b) Octahedral shape
   (c) Steric hindrance protecting the Sulphur atom
   (d) Presence of d-orbitals in Sulphur

8. Ozone (O₃) acts as a powerful oxidizing agent because:
   (a) It is thermodynamically unstable and readily decomposes to give O₂ and nascent oxygen [O]
   (b) It contains coordinate bonds
   (c) It is diamagnetic
   (d) It has a bent structure

9. Which of the following interhalogen compounds has a bent T-shape?
   (a) IF₇
   (b) IF₅
   (c) ClF
   (d) BrF₃

10. The Haber's process for the manufacture of ammonia uses the following conditions:
    (a) High temperature (~700K), High pressure (~200 atm), FeO catalyst
    (b) Low temperature, Low pressure, V₂O₅ catalyst
    (c) High temperature (~500K), Low pressure (~1 atm), Pt catalyst
    (d) Low temperature, High pressure (~500 atm), Ni catalyst

Answer Key:

1. (b) H₂S (H₂O has abnormally high BP due to H-bonding. Down the group, BP increases due to van der Waals forces, so H₂S is lowest among H₂S, H₂Se, H₂Te).
2. (c) Inert pair effect
3. (c) Square planar
4. (b) H₃PO₃ (It has two P-OH bonds)
5. (b) Cl > F > Br > I
6. (c) NO₂
7. (c) Steric hindrance protecting the Sulphur atom (Kinetically inert)
8. (a) It is thermodynamically unstable and readily decomposes to give O₂ and nascent oxygen [O]
9. (d) BrF₃ (also ClF₃)
10. (a) High temperature (~700K), High pressure (~200 atm), FeO catalyst (with promoters K₂O, Al₂O₃)

Study these notes thoroughly. Practice drawing structures and writing balanced equations. Good luck with your preparation!