Class 12 Chemistry Notes Chapter 7 (The p- Block Elements) – Examplar Problems Book

Alright class, let's dive deep into Chapter 7, 'The p-Block Elements'. This is a very important chapter for your competitive exams, covering Groups 15 to 18. Pay close attention to the trends, structures, and key reactions.

The p-Block Elements (Groups 13-18)

Elements in which the last electron enters any of the three p-orbitals of their outermost shell are called p-block elements.

• General Electronic Configuration: ns²np¹⁻⁶
• Class 12 syllabus focuses on Groups 15, 16, 17, and 18.
• These groups show great variation: contain non-metals, metalloids, and metals.

Group 15 Elements (Nitrogen Family)

• Elements: Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb), Bismuth (Bi).
• Electronic Configuration: ns²np³ (half-filled p-orbitals, hence extra stability).
• Occurrence: N₂ (78% of air), Minerals (Nitrates like NaNO₃, KNO₃), Proteins. P (Apatite family, e.g., Ca₉(PO₄)₆.CaX₂, X=F, Cl or OH), Bones, Nucleic acids. As, Sb, Bi mainly as sulfide minerals.

Trends in Properties (Group 15):

1. Atomic/Ionic Radii: Increase down the group due to the addition of a new shell. Covalent radius increases significantly from N to P, but only a small increase from As to Bi (due to poor shielding by d and f orbitals).
2. Ionization Enthalpy (IE): Decreases down the group due to increasing atomic size. Group 15 elements have higher IE than Group 14 due to smaller size and stable half-filled configuration. IE₁ > IE₂ > IE₃.
3. Electronegativity: Decreases down the group due to increasing atomic size.
4. Physical Properties: All elements are polyatomic. N₂ is a diatomic gas. Others are solids. Metallic character increases down the group (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.
5. Oxidation States: Common: -3, +3, +5.
   • Tendency to show -3 state decreases down the group (due to increasing size and metallic character).
   • Stability of +5 state decreases down the group, while +3 state increases. This is due to the Inert Pair Effect (reluctance of ns² electrons to participate in bonding due to poor shielding by inner d and f electrons). Bi predominantly shows +3.
   • Nitrogen exhibits a wide range of oxidation states (-3, -2, -1, 0, +1, +2, +3, +4, +5) due to its ability to form multiple bonds with oxygen.
6. Anomalous Behaviour of Nitrogen: Due to its small size, high electronegativity, high ionization enthalpy, and absence of d-orbitals.
   • Forms pπ–pπ multiple bonds (N≡N triple bond is very strong, making N₂ inert). Heavier elements don't form pπ–pπ bonds effectively.
   • Cannot form dπ–pπ bonds.
   • Maximum covalency is 4 (using s and p orbitals). Heavier elements can expand covalency using d-orbitals (e.g., PCl₅).

Important Compounds (Group 15):

1. Dinitrogen (N₂):

   • Preparation: Lab: NH₄Cl(aq) + NaNO₂(aq) → N₂(g) + 2H₂O(l) + NaCl(aq). Pure N₂: Ba(N₃)₂(s) → Ba(s) + 3N₂(g) (Thermal decomposition of azides).
   • Properties: Colourless, odourless, tasteless, non-toxic gas. Very inert at room temperature due to high N≡N bond enthalpy (941.4 kJ/mol). Reacts at high temperatures.
   • Uses: Manufacturing NH₃, HNO₃; Inert atmosphere; Cryosurgery; Refrigerant.

2. Ammonia (NH₃):

   • Preparation: Haber's Process: N₂(g) + 3H₂(g) ⇌ 2NH₃(g); ΔfH° = –46.1 kJ/mol. (High pressure ~200 atm, Temp ~700 K, Catalyst: Iron oxide with K₂O and Al₂O₃ promoters).
   • Structure: Trigonal pyramidal, sp³ hybridized, one lone pair.
   • Properties: Colourless gas, pungent smell. Highly soluble in water (forms NH₄OH, weak base). Acts as a Lewis base (donates lone pair), forms coordination compounds.
   • Uses: Production of fertilizers (urea, ammonium nitrate), HNO₃, refrigerant.

3. Oxides of Nitrogen: (Know structures, oxidation states, preparation methods)

   • N₂O (+1), NO (+2, paramagnetic), N₂O₃ (+3), NO₂ (+4, paramagnetic, dimerises to N₂O₄), N₂O₄ (+4), N₂O₅ (+5).

4. Nitric Acid (HNO₃):

   • Preparation: Ostwald's Process:
     1. Catalytic oxidation of NH₃: 4NH₃ + 5O₂ --(Pt/Rh gauge, 500K, 9 bar)--> 4NO + 6H₂O
     2. Oxidation of NO: 2NO + O₂ ⇌ 2NO₂
     3. Absorption in water: 3NO₂ + H₂O → 2HNO₃ + NO (NO is recycled)
   • Properties: Colourless liquid. Strong oxidizing agent (reacts with most metals, except noble metals like Au, Pt). Oxidizing action depends on concentration, temperature, and nature of the material. Forms passive layer on Cr, Al. Brown Ring Test for nitrates: Fe²⁺ + NO₃⁻ + 4H⁺ → Fe³⁺ + NO + 2H₂O; [Fe(H₂O)₆]²⁺ + NO → [Fe(H₂O)₅(NO)]²⁺ (Brown complex).
   • Uses: Manufacturing fertilizers, explosives (TNT), dyes, drugs.

5. Phosphorus (P):

   • Allotropes: White (or Yellow), Red, Black.
     • White P: Discrete tetrahedral P₄ units. Highly reactive, poisonous, chemiluminescent. Soluble in CS₂. Angle strain (60°).
     • Red P: Polymeric structure (P₄ units linked). Less reactive, non-poisonous. Insoluble in CS₂.
     • Black P: Most stable. Two forms (α and β). Layered structure.

6. Phosphine (PH₃):

   • Preparation: Ca₃P₂ + 6H₂O → 3Ca(OH)₂ + 2PH₃. Lab: P₄ + 3NaOH + 3H₂O → PH₃ + 3NaH₂PO₂ (Sodium hypophosphite).
   • Properties: Colourless gas, rotten fish smell. Poisonous. Weakly basic (less than NH₃). Spontaneously flammable due to P₂H₄ or P₄ vapour impurities.
   • Uses: Holme's signals (containers of CaC₂ and Ca₃P₂ thrown in sea produce gases that ignite).

7. Phosphorus Halides:

   • PCl₃: Preparation: P₄ + 6Cl₂ → 4PCl₃. Structure: Pyramidal (sp³). Hydrolysis: PCl₃ + 3H₂O → H₃PO₃ + 3HCl.
   • PCl₅: Preparation: P₄ + 10Cl₂ → 4PCl₅. Structure: Trigonal bipyramidal (sp³d) in gas/liquid. Exists as [PCl₄]⁺[PCl₆]⁻ (ionic) in solid state. Hydrolysis: PCl₅ + H₂O → POCl₃ + 2HCl; POCl₃ + 3H₂O → H₃PO₄ + 3HCl.

8. Oxoacids of Phosphorus: (Focus on structure, basicity, and reducing character)

   • Hypophosphorous acid (H₃PO₂): Monobasic (one P-OH bond), Strong reducing agent (two P-H bonds).
   • Orthophosphorous acid (H₃PO₃): Dibasic (two P-OH bonds), Reducing agent (one P-H bond). Disproportionates on heating.
   • Orthophosphoric acid (H₃PO₄): Tribasic (three P-OH bonds). Non-reducing.
   • Pyrophosphoric acid (H₄P₂O₇): Tetrabasic. Formed by heating H₃PO₄.
   • Metaphosphoric acid ((HPO₃)n): Cyclic or polymeric.

Group 16 Elements (Chalcogens)

• Elements: Oxygen (O), Sulfur (S), Selenium (Se), Tellurium (Te), Polonium (Po) (Radioactive).
• Electronic Configuration: ns²np⁴.
• Occurrence: O (Air, Earth's crust). S (Sulfides - ZnS, PbS; Sulfates - CaSO₄·2H₂O). Se, Te as selenides, tellurides in sulfide ores. Po (decay product of Th, U).

Trends in Properties (Group 16):

1. Atomic/Ionic Radii: Increase down the group.
2. Ionization Enthalpy: Decrease down the group. Lower than Group 15 (due to stable half-filled config in Group 15).
3. Electron Gain Enthalpy: Oxygen has less negative EGE than Sulfur (due to compact size of O). Down the group, value becomes less negative. S > Se > Te > Po > O.
4. Electronegativity: Decrease down the group. Oxygen is the second most electronegative element.
5. Physical Properties: O₂ is gas, others are solids. Metallic character increases down the group (O, S - non-metals; Se, Te - metalloids; Po - metal). Melting/Boiling points increase down the group. Large difference between O and S due to atomicity (O₂, S₈).
6. Oxidation States: Common: -2, +2, +4, +6.
   • Oxygen mostly shows -2 (except OF₂ (+2), H₂O₂ (-1), KO₂ (-1/2)).
   • Stability of -2 state decreases down the group.
   • Stability of +6 state decreases, +4 increases down the group (Inert Pair Effect). S, Se, Te show +4, +6 due to d-orbitals.
7. Anomalous Behaviour of Oxygen: Due to small size, high electronegativity, absence of d-orbitals.
   • Forms strong pπ–pπ bonds (O=O).
   • H-bonding in H₂O leads to liquid state (unlike H₂S gas).
   • Maximum covalency usually 2 (can extend to 4 in H₃O⁺), while others show up to 6 (e.g., SF₆).

Important Compounds (Group 16):

1. Dioxygen (O₂):

   • Preparation: Lab: KClO₃ --(MnO₂)--> 2KCl + 3O₂. Thermal decomposition of oxides (Ag₂O, HgO, PbO₂). H₂O₂ --(MnO₂)--> 2H₂O + O₂.
   • Properties: Colourless, odourless gas. Paramagnetic (due to unpaired electrons in π* antibonding MOs). Reacts with most metals and non-metals.
   • Uses: Respiration, Combustion, Oxy-acetylene welding, Steel production.

2. Ozone (O₃):

   • Preparation: Passing silent electric discharge through pure, cold, dry oxygen (prevents decomposition back to O₂). 3O₂ ⇌ 2O₃; ΔH = +142 kJ/mol (endothermic).
   • Structure: Angular/Bent molecule. Resonance hybrid. O-O bond lengths are equal.
   • Properties: Pale blue gas, fishy smell. Thermodynamically unstable compared to O₂. Powerful oxidizing agent (stronger than O₂). O₃ → O₂ + [O] (nascent oxygen). Oxidises PbS to PbSO₄, KI to I₂ (used for quantitative estimation). Depleted by CFCs in the stratosphere.
   • Uses: Germicide, disinfectant, sterilizing water, bleaching oils, flour, starch.

3. Sulfur (S):

   • Allotropes: Yellow Rhombic (α-sulfur) and Monoclinic (β-sulfur). Both have puckered S₈ rings (crown shape). Rhombic stable below 369 K, Monoclinic above 369 K (Transition Temperature). S₂ exists at high temp (~1000 K), paramagnetic like O₂.

4. Sulfur Dioxide (SO₂):

   • Preparation: Burning S in air: S + O₂ → SO₂. Roasting sulfide ores: 4FeS₂ + 11O₂ → 2Fe₂O₃ + 8SO₂. Lab: Na₂SO₃ + H₂SO₄(dil) → Na₂SO₄ + H₂O + SO₂.
   • Structure: Angular/Bent (sp²).
   • Properties: Colourless gas, pungent smell. Highly soluble in water (forms sulfurous acid, H₂SO₃). Acts as a reducing agent (e.g., decolorizes acidified KMnO₄, reduces Fe³⁺ to Fe²⁺). Also acts as an oxidizing agent (e.g., with H₂S). Bleaching action (temporary, due to reduction) on delicate articles (wool, silk).
   • Uses: Refining petroleum/sugar, Bleaching, Disinfectant, Manufacturing H₂SO₄, NaHSO₃.

5. Oxoacids of Sulfur: (Know structures)

   • Sulfurous acid (H₂SO₃) - Dibasic
   • Sulfuric acid (H₂SO₄) - Dibasic
   • Peroxodisulfuric acid (Marshall's acid, H₂S₂O₈) - Contains peroxide linkage (-O-O-)
   • Pyrosulfuric acid (Oleum, H₂S₂O₇) - Contains S-O-S linkage.

6. Sulfuric Acid (H₂SO₄): "King of Chemicals"

   • Preparation: Contact Process:
     1. Burning S or roasting ores to get SO₂.
     2. Catalytic oxidation of SO₂ to SO₃: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g); ΔH = -196.6 kJ/mol (Catalyst: V₂O₅, Temp: ~720 K, Pressure: ~2 bar). Exothermic, reversible - Le Chatelier's principle applied.
     3. Absorption of SO₃ in conc. H₂SO₄ to form Oleum (H₂S₂O₇). SO₃ + H₂SO₄ → H₂S₂O₇.
     4. Dilution of Oleum with water: H₂S₂O₇ + H₂O → 2H₂SO₄.
   • Properties: Colourless, dense, oily liquid. High boiling point, viscous (due to H-bonding). Strong dibasic acid. Low volatility. Strong dehydrating agent (chars carbohydrates C₁₂H₂₂O₁₁ --(H₂SO₄)--> 12C + 11H₂O). Moderate strength oxidizing agent (oxidizes metals, non-metals; products depend on temp & conc.).
   • Uses: Fertilizer production (ammonium sulfate), manufacturing chemicals, petroleum refining, detergents, metallurgy, batteries.

Group 17 Elements (Halogens)

• Elements: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At) (Radioactive).
• Electronic Configuration: ns²np⁵.
• Occurrence: Highly reactive, exist only as compounds (halides). Fluorspar (CaF₂), Cryolite (Na₃AlF₆), NaCl (sea water), Carnallite (KCl.MgCl₂.6H₂O).

Trends in Properties (Group 17):

1. Atomic/Ionic Radii: Smallest in respective periods. Increase down the group.
2. Ionization Enthalpy: Very high. Decrease down the group.
3. Electron Gain Enthalpy (EGE): Maximum negative EGE in the periodic table. Anomaly: EGE of Cl is more negative than F. (Reason: Small size of F leads to strong inter-electronic repulsions in the compact 2p subshell, making addition of electron less favourable than in Cl). Order: Cl > F > Br > I.
4. Electronegativity: Highest in respective periods. Decrease down the group. F is the most electronegative element.
5. Physical Properties: Exist as diatomic molecules (X₂). F₂, Cl₂ (gases), Br₂ (liquid), I₂ (solid). Colour arises from absorption of radiation in visible region, exciting outer electrons. Colour deepens down the group (F - pale yellow, Cl - greenish yellow, Br - red-brown, I - violet solid/gas). Low M.P/B.P, increase down the group due to stronger van der Waals forces.
6. Bond Dissociation Enthalpy (BDE): Anomaly: Cl₂ > Br₂ > F₂ > I₂. (Reason: Low BDE of F₂ is due to large electron-electron repulsion among lone pairs in the small F₂ molecule).
7. Oxidation States: Usually -1. F shows only -1 (most electronegative). Cl, Br, I also show positive oxidation states (+1, +3, +5, +7) in compounds with more electronegative elements (O, F) and in oxoacids/interhalogens, due to available d-orbitals.
8. Oxidizing Power: Strong oxidizing agents. Oxidizing power decreases down the group (F₂ > Cl₂ > Br₂ > I₂). F₂ oxidizes water. Cl₂, Br₂ can oxidize ions lower down. (e.g., Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂).
9. Anomalous Behaviour of Fluorine: Due to small size, highest electronegativity, low F-F BDE, absence of d-orbitals. Shows only -1 oxidation state. HF is liquid (H-bonding), others are gases.

Important Compounds (Group 17):

1. Chlorine (Cl₂):

   • Preparation: MnO₂ + 4HCl(conc) → MnCl₂ + Cl₂ + 2H₂O. KMnO₄ + HCl. Deacon's process: 4HCl + O₂ --(CuCl₂)--> 2Cl₂ + 2H₂O. Electrolysis of brine (NaCl solution).
   • Properties: Greenish-yellow gas, pungent suffocating odour. Soluble in water (forms HCl and HOCl - hypochlorous acid). HOCl is unstable, decomposes to HCl and nascent oxygen [O]. This [O] is responsible for oxidizing and bleaching properties (permanent bleaching by oxidation). Reacts with metals, non-metals. Reacts with NH₃ (excess NH₃ gives N₂, excess Cl₂ gives NCl₃ explosive). Reacts with NaOH (cold/dilute gives NaCl + NaOCl; hot/conc gives NaCl + NaClO₃).
   • Uses: Bleaching wood pulp, cotton; Sterilizing water; Extraction of Au, Pt; Manufacturing dyes, drugs, DDT, CFCs, HCl.

2. Hydrogen Chloride (HCl):

   • Preparation: Lab: NaCl + H₂SO₄(conc) --(420K)--> NaHSO₄ + HCl. NaHSO₄ + NaCl --(823K)--> Na₂SO₄ + HCl.
   • Properties: Colourless, pungent gas. Easily liquefied. Extremely soluble in water (forms hydrochloric acid). Aqueous solution is strong acid. Reacts with metals, bases. Forms aqua regia (3 parts conc HCl + 1 part conc HNO₃) which dissolves noble metals (Au, Pt).
   • Uses: Manufacturing Cl₂, NH₄Cl, glucose; Extracting glue from bones; Refining sugar; Medicine; Lab reagent.

3. Oxoacids of Halogens: (Focus on structure, acidic strength, oxidizing power)

   • Hypohalous acid (HOX), Halous acid (HXO₂ - only for Cl), Halic acid (HXO₃), Perhalic acid (HXO₄).
   • Acidic strength increases with increasing oxidation state of halogen: HOX < HXO₂ < HXO₃ < HXO₄. (Reason: Stability of conjugate base increases).
   • Oxidizing power decreases with increasing oxidation state: HOCl > HClO₂ > HClO₃ > HClO₄.
   • Structures based on VSEPR theory.

4. Interhalogen Compounds: Compounds formed between two different halogens (X-X').

   • Types: XX', XX'₃, XX'₅, XX'₇ (where X is larger/less electronegative halogen).
   • Preparation: Direct combination (e.g., Cl₂ + F₂(equal vol) → 2ClF; I₂ + 3Cl₂(excess) → 2ICl₃).
   • Properties: Covalent molecules, diamagnetic. More reactive than constituent halogens (except F₂) because X-X' bond is weaker than X-X or X'-X' bond. Undergo hydrolysis (products depend on specific compound, generally halide ion of smaller halogen and oxyhalide ion of larger halogen).
   • Structures: XX' (Linear), XX'₃ (Bent T-shape, e.g., ClF₃), XX'₅ (Square pyramidal, e.g., BrF₅), XX'₇ (Pentagonal bipyramidal, e.g., IF₇). Based on VSEPR theory (consider lone pairs).
   • Uses: Non-aqueous solvents, Fluorinating agents (e.g., ClF₃, BrF₃ used in UF₆ production).

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). Completely filled valence shell.
• Occurrence: All except Rn occur in atmosphere (Ar is major component ~1%). He found in natural gas. Rn is decay product of Radium.

Trends in Properties (Group 18):

1. Atomic Radii: Increase down the group. Largest in their respective periods (van der Waals radii considered).
2. Ionization Enthalpy: Extremely high due to stable configuration. Decreases down the group.
3. Electron Gain Enthalpy: Large positive values (energy has to be supplied to add an electron).
4. Physical Properties: Monoatomic gases. Colourless, odourless, tasteless. Sparingly soluble in water. Very low melting and boiling points (weak van der Waals/dispersion forces). M.P/B.P increase down the group with increasing atomic size and magnitude of van der Waals forces. He has the lowest B.P of any known substance (4.2 K).

Chemical Properties (Group 18):

• Generally inert due to stable configuration, high IE, positive EGE.
• Discovery of Compounds: Neil Bartlett (1962) observed PtF₆ oxidizing O₂ to O₂⁺[PtF₆]⁻. Since IE₁ of O₂ (1175 kJ/mol) was close to IE₁ of Xe (1170 kJ/mol), he reacted Xe with PtF₆ and prepared the first noble gas compound, Xe⁺[PtF₆]⁻ (red solid).
• Subsequently, many compounds of Xe (mainly with highly electronegative F and O) were synthesized. Kr compounds are fewer (e.g., KrF₂). Rn compounds difficult to study (radioactive). No true compounds of He, Ne, Ar known.

Xenon Compounds:

1. Xenon Fluorides:

   • Xe(g) + F₂(g) --(673K, 1bar, Ni tube)--> XeF₂(s) (Xe in excess)
   • Xe(g) + 2F₂(g) --(873K, 7bar)--> XeF₄(s) (Xe:F₂ = 1:5)
   • Xe(g) + 3F₂(g) --(573K, 60-70bar)--> XeF₆(s) (Xe:F₂ = 1:20)
   • Properties: Colourless crystalline solids. Powerful fluorinating agents. Readily hydrolysed.
   • Hydrolysis:
     • 2XeF₂(s) + 2H₂O(l) → 2Xe(g) + 4HF(aq) + O₂(g) (Complete)
     • 6XeF₄(s) + 12H₂O(l) → 4Xe(g) + 2XeO₃(s) + 24HF(aq) + 3O₂(g) (Complete, disproportionation)
     • XeF₆(s) + 3H₂O(l) → XeO₃(s) + 6HF(aq) (Complete)
     • XeF₆ + H₂O → XeOF₄ + 2HF (Partial)
     • XeF₆ + 2H₂O → XeO₂F₂ + 4HF (Partial)
   • Structures: XeF₂ (Linear, sp³d), XeF₄ (Square planar, sp³d²), XeF₆ (Distorted octahedral, sp³d³). Based on VSEPR (Xe has lone pairs).

2. Xenon Oxides & Oxyfluorides:

   • XeO₃: Prepared by hydrolysis of XeF₄ or XeF₆. Colourless explosive solid. Pyramidal structure (sp³).
   • XeOF₄: Partial hydrolysis of XeF₆. Colourless volatile liquid. Square pyramidal structure (sp³d²).
   • XeO₂F₂: Partial hydrolysis of XeF₆. Trigonal bipyramidal structure (sp³d, see-saw like geometry if lone pair considered in equatorial).

Uses of Noble Gases:

• He: Non-inflammable, light gas - filling balloons, airships. Liquid He - cryogenics. Used in gas-cooled nuclear reactors. Oxygen mixture for deep-sea divers (less soluble in blood than N₂). NMR spectrometers.
• Ne: Discharge tubes, fluorescent bulbs ("neon signs"). Botanical gardens, greenhouses.
• Ar: Inert atmosphere in welding, metallurgy. Filling electric bulbs (reduces filament evaporation).
• Kr, Xe: Special purpose light bulbs, high-intensity lamps.
• Rn: Radiotherapy (cancer treatment).

Multiple Choice Questions (MCQs):

1. Which of the following elements exhibits the most pronounced inert pair effect?
   (a) N
   (b) P
   (c) Sb
   (d) Bi
   Answer: (d) Bi (Inert pair effect increases down the group, making the lower oxidation state, +3 for Bi, more stable than +5).

2. The geometry of the ammonia molecule (NH₃) can be best described as:
   (a) Trigonal planar
   (b) Tetrahedral
   (c) Trigonal pyramidal
   (d) Square planar
   Answer: (c) Trigonal pyramidal (sp³ hybridization with one lone pair leads to pyramidal geometry).

3. Which process is used for the commercial production of nitric acid?
   (a) Haber's Process
   (b) Contact Process
   (c) Ostwald's Process
   (d) Deacon's Process
   Answer: (c) Ostwald's Process (Involves catalytic oxidation of ammonia).

4. Oxygen molecule (O₂) is paramagnetic. This is best explained by:
   (a) VSEPR Theory
   (b) Valence Bond Theory
   (c) Molecular Orbital Theory
   (d) Lewis Structure
   Answer: (c) Molecular Orbital Theory (MOT explains the presence of two unpaired electrons in the π* antibonding orbitals).

5. Which of the following oxoacids of sulfur contains an S-O-S linkage?
   (a) H₂SO₃ (Sulfurous acid)
   (b) H₂SO₄ (Sulfuric acid)
   (c) H₂S₂O₇ (Pyrosulfuric acid/Oleum)
   (d) H₂S₂O₈ (Peroxodisulfuric acid)
   Answer: (c) H₂S₂O₇ (Oleum is formed by dissolving SO₃ in H₂SO₄, creating an S-O-S bridge).

6. The correct order of decreasing negative electron gain enthalpy among halogens is:
   (a) F > Cl > Br > I
   (b) Cl > F > Br > I
   (c) Br > Cl > F > I
   (d) I > Br > Cl > F
   Answer: (b) Cl > F > Br > I (Chlorine has the most negative EGE due to optimal size; Fluorine's EGE is lower than expected due to small size and interelectronic repulsion).

7. Bleaching action of chlorine is due to:
   (a) Reduction
   (b) Oxidation
   (c) Chlorination
   (d) Hydrogenation
   Answer: (b) Oxidation (Cl₂ reacts with water to form HOCl, which releases nascent oxygen [O], a strong oxidizing agent responsible for permanent bleaching).

8. Which interhalogen compound has a square pyramidal geometry?
   (a) ClF₃
   (b) IF₇
   (c) BrF₅
   (d) XeF₄
   Answer: (c) BrF₅ (Central Br atom is sp³d² hybridized with one lone pair, leading to square pyramidal geometry. XeF₄ is square planar).

9. The first noble gas compound prepared by Neil Bartlett involved:
   (a) Argon
   (b) Krypton
   (c) Xenon
   (d) Helium
   Answer: (c) Xenon (He prepared Xe⁺[PtF₆]⁻ based on the similarity in ionization enthalpy between O₂ and Xe).

10. Which noble gas is used in cryogenic applications due to its extremely low boiling point?
    (a) Neon
    (b) Argon
    (c) Krypton
    (d) Helium
    Answer: (d) Helium (Liquid helium has the lowest boiling point (4.2 K) and is extensively used as a cryogenic agent).

Study these notes thoroughly, focusing on understanding the concepts rather than just memorizing facts. Good luck with your preparation!