Unit 7 – The p-Block Elements Class 12 Notes | Group 15, 16, 17, 18 Elements Explained | CBSE Chemistry 2025

 Unit 7: The p-Block Elements 

7.1 Group 15 Elements

7.2 Dinitrogen

7.3 Ammonia

7.4 Oxides of Nitrogen

7.5 Nitric Acid

7.6 Phosphorus – Allotropic Forms

7.7 Phosphine

7.8 Phosphorus Halides

7.9 Oxoacids of Phosphorus

7.10 Group 16 Elements

7.11 Dioxygen

7.12 Simple Oxides

7.13 Ozone

7.14 Sulphur – Allotropic Forms

7.15 Sulphur Dioxide

7.16 Oxoacids of Sulphur

7.17 Sulphuric Acid

7.18 Group 17 Elements

7.19 Chlorine

7.20 Hydrogen Chloride

7.21 Oxoacids of Halogens

7.22 Interhalogen Compounds

7.23 Group 18 Elements

I. General Properties and Trends in p-Block Elements

• Affinity for Hydrogen and Bond Dissociation Enthalpy
  • The affinity for hydrogen decreases as you go down the group from fluorine to iodine within the halogens ``.
  • Consequently, the highest bond dissociation enthalpy among halogen acids (HF, HCl, HBr, HI) is found in HF ``. This implies a stronger bond that is harder to break.
• Reducing Behavior and Bond Dissociation Enthalpy
  • The strength of a reducing agent among E—H compounds (where E is an element like N, P, As, Sb) is inversely related to the bond dissociation enthalpy of the E—H bond ``.
  • A lower bond dissociation enthalpy indicates a weaker bond, making it easier to release hydrogen and act as a reducing agent ``.
  • For the compounds NH₃, PH₃, AsH₃, and SbH₃, the bond dissociation enthalpy (∆ diss (E—H)) decreases from NH₃ (389 kJ mol⁻¹) to SbH₃ (255 kJ mol⁻¹) ``.
  • Therefore, SbH₃ will act as the strongest reducing agent ``.
• Electron Gain Enthalpy and Thermal Stability
  • The order of more negative electron gain enthalpy for O, S, Cl, F is S < O < Cl < F ``. (Note: The source presents this as a statement to be evaluated, and implies it is not correct in general, but the specific ordering provided is S < O < Cl < F which is incorrect because Chlorine has the highest electron gain enthalpy, followed by Fluorine, then Oxygen, then Sulfur. However, the question 34 part (iii) states S < O < Cl < F for More negative electron gain enthalpy and is marked as incorrect, implying the correct understanding is that Cl and F are higher, and the order given is wrong).
  • Thermal stability of hydrides H₂O > H₂S > H₂Se > H₂Te is a correct order ``. This generally decreases down the group.
• Acid Strength of Oxides
  • The given order for acid strength As₂O₃ < SiO₂ < P₂O₃ < SO₂ is stated as correct in the source ``. Acid strength of oxides generally increases across a period and down a group for non-metals.

II. Properties and Reactions of Specific p-Block Elements and Compounds

• Reactions with Sulfuric Acid (H₂SO₄)
  • Concentrated H₂SO₄ ``
    • When concentrated H₂SO₄ is added to a chloride salt, colourless fumes are evolved ``.
    • When added to an iodide salt, violet fumes are observed because HI is oxidised to I₂ by H₂SO₄ ``.
    • Hot concentrated H₂SO₄ acts as a moderately strong oxidising agent ``.
    • It oxidizes Copper (Cu) to CuSO₄ + SO₂ + 2H₂O ``.
    • It oxidizes Sulphur (S) to two gaseous products ``.
    • It oxidizes Carbon (C) into two gaseous products ``.
    • It oxidizes Zinc (Zn) ``.
    • Reactions where H₂SO₄ acts as an oxidizing reagent include: 2HI + H₂SO₄ → I₂ + SO₂ + 2H₂O and Cu + 2H₂SO₄ → CuSO₄ + SO₂ + 2H₂O ``.
    • In the preparation of H₂SO₄ by the Contact Process, SO₃ is not absorbed directly in water to form H₂SO₄ because it forms an acid fog, which is difficult to condense ``.
• Nitrogen and its Compounds
  • Allotropy: Nitrogen does not show allotropy ``.
  • Covalency: The maximum covalency of nitrogen is four ``.
  • Bond Strength: The single N–N bond is weaker than the single P–P bond . (The statement "Single N–N bond is stronger than the single P–P bond" is explicitly stated as wrong in the source).
  • Reactivity: N₂ is less reactive than P₄ . The reason provided that "Nitrogen has more electron gain enthalpy than phosphorus" is given as an incorrect explanation.
  • Nitric Oxide (NO):
    • NO₂ is paramagnetic in nature ``.
    • Nitric oxide (NO) is paramagnetic in the gaseous state because it exists as a monomer with one unpaired electron ``.
    • On cooling, it dimerises to N₂O₄ in the solid state, resulting in no unpaired electrons, thus the solid form is diamagnetic ``.
  • Oxides of Nitrogen:
    • On heating lead nitrate, the oxides formed are NO₂ and PbO ``.
    • Nitric acid (HNO₃) reacts with P₄O₁₀ to form 4HPO₃ + 2N₂O₅ ``.
    • Effect of concentration of nitric acid on oxidation product: Dilute and concentrated nitric acid give different oxidation products when reacting with copper metal ``:
      • 3Cu + 8HNO₃ (dil.) → 3Cu(NO₃)₂ + 2NO + 4H₂O ``
      • Cu + 4HNO₃ (Conc.) → 3Cu(NO₃)₂ + 2NO + 2H₂O `` (Note: there is a typo in the second reaction's product, it should produce NO2. However, the source provided the reaction as given).
    • Nitric Acid and Iron: HNO₃ makes iron passive . The reason given is that it forms a protective layer of **ferric nitrate** on the surface of iron, which is stated as a wrong reason.
  • Ammonia (NH₃):
    • On heating ammonium dichromate, N₂ gas is produced ``.
    • Catalytic oxidation of NH₃ by atmospheric oxygen: 4NH₃ + 5O₂ (from air) --(Pt/Rh gauge catalyst, 500K, 9 bar)→ 4NO + 6H₂O ``.
    • NH₃ dissolves in water because it forms hydrogen bonds with water ``.
    • It is more basic than the gas (phosphorine, PH₃) produced from white phosphorus and concentrated NaOH solution . (The statement that it is *less* basic than the gas is incorrect).
    • In the Haber's process for NH₃ preparation, iron powder along with Al₂O₃ and K₂O is used as a catalyst ``.
  • Oxoacids of Nitrogen: Three oxoacids are HNO₂, HNO₃, and H₂N₂O₂ (Hyponitrous acid) ``.
    • Disproportionation reaction of HNO₂ (Nitrogen in +3 oxidation state): 3HNO₂ → HNO₃ + H₂O + 2NO ``.
• Phosphorus and its Compounds
  • pπ–dπ bonding: Phosphorus can be involved in pπ–dπ bonding ``. Carbon, Nitrogen, and Boron generally do not, or have very limited d-orbital participation compared to elements in lower periods.
  • White Phosphorus (P₄):
    • On heating with concentrated NaOH solution in an inert atmosphere of CO₂, white phosphorus gives a gas which is highly poisonous and has a smell like rotten fish ``.
    • P₄ molecule of white phosphorus has six P–P single bonds and four lone pairs of electrons . It does not have three P-P single bonds.
    • In white phosphorus, there is angular strain in P₄ molecules because bond angles are only 60°, making white phosphorus very reactive ``.
    • White phosphorus is a discrete tetrahedral molecule ``.
  • Red Phosphorus: Has a polymeric structure where P₄ tetrahedra are linked through P—P bonds to form chains . It is **less reactive** than white phosphorus.
  • Phosphorus Halides:
    • On reaction with Cl₂, phosphorus forms two halides:
      • A is PCl₅ (yellowish-white powder) . Reaction: `P₄ + 10Cl₂ → 4PCl₅`. Its hydrolysis product is H₃PO₄ + 5HCl from PCl₅ + 4H₂O ``.
      • B is PCl₃ (colourless oily liquid) . Reaction: `P₄ + 6Cl₂ → 4PCl₃`. Its hydrolysis product is H₃PO₃ + 3HCl from PCl₃ + 3H₂O ``.
    • In PCl₅, phosphorus is in sp³d hybridised state, but all its five bonds are not equivalent . This is due to its **trigonal bipyramidal geometry**, where axial and equatorial bonds differ in length and strength.
    • In the solid state, PCl₅ is an ionic solid with [PCl₄]⁺ tetrahedral and [PCl₆]⁻ octahedral ions . (The statement "PCl is ionic in solid state in which cation is tetrahedral and anion is octahedral" is correct, assuming PCl refers to PCl₅ in solid state).
    • When PCl₅ reacts with finely divided silver on heating, AgCl (a white silver salt) is obtained, which dissolves upon adding excess aqueous NH₃ solution to form a soluble complex ``:
      • PCl₅ + 2Ag → 2AgCl + PCl₃ ``
      • AgCl + 2NH₃(aq) → [Ag(NH₃)₂]⁺Cl⁻ (soluble complex) ``
  • Oxoacids of Phosphorus:
    • H₃PO₄ (Phosphoric acid) forms three series of salts ``. This means it is a triprotic acid.
    • H₃PO₂ (Phosphinic acid/Hypophosphorous acid) has strong reducing behaviour due to the presence of one –OH group and two P–H bonds . Its structure features one P=O, one P-OH, and two P-H bonds.
    • A reaction showing its reducing behaviour: 4AgNO₃ + 2H₂O + H₃PO₂ → 4Ag + 4HNO₃ + H₃PO₄ ``.
    • Pyrophosphoric acid (H₄P₂O₇) has a specific structure ``.
• Sulfur and its Compounds
  • Peroxoacids of Sulfur: H₂SO₅ (Peroxomonosulphuric acid) and H₂S₂O₈ (Peroxodisulphuric acid) are peroxoacids of sulphur ``.
    • In peroxosulphuric acid (H₂SO₅), sulphur is in +6 oxidation state ``.
  • SO₂ gas:
    • It acts as a bleaching agent in moist conditions ``.
    • Its dilute solution is used as a disinfectant ``.
    • Its molecule does not have linear geometry ``.
    • It cannot be prepared by the reaction of dilute H₂SO₄ with metal sulphide ``.
  • SF₆ vs SCl₆: SF₆ is known but SCl₆ is not . This is because the **small size of fluorine allows six F⁻ ions to be accommodated around sulphur**, while the larger size of chloride ions would lead to **interionic repulsion**.
  • Sulphur Allotropes: Both rhombic and monoclinic sulphur exist as S₈ . Unlike oxygen, which forms pπ–pπ multiple bonds due to its small size, pπ–pπ bonding is **not possible in sulphur**.
• Halogens (Chlorine, Iodine) and their Compounds
  • Reaction of Chlorine with NaOH: If chlorine gas is passed through hot NaOH solution, two changes are observed in the oxidation number of chlorine during the reaction: 0 to +5 and 0 to –1 ``. This is a disproportionation reaction.
  • Oxidising Power: The order F₂ > Cl₂ > Br₂ > I₂ correctly represents the oxidising power ``.
  • Bond Dissociation Enthalpy: The order F₂ > Cl₂ > Br₂ > I₂ for bond dissociation enthalpy is incorrect ``.
  • Hydrogen-Halogen Bond Strength: The order HI < HBr < HCl < HF for hydrogen-halogen bond strength is correct ``.
  • Electron Gain Enthalpy: Chlorine (Cl₂) has the highest electron gain enthalpy ``.
  • ClO₃⁻ and SO₃²⁻ are an isoelectronic and isostructural pair . (Note: Other options like CO₃²⁻, NO₃⁻ are also isoelectronic and isostructural).
  • Stability of Oxoacids of Chlorine: The stability of oxoacids of chlorine increases in the order HClO < HClO₂ < HClO₃ < HClO₄ . This is because the **dispersal of negative charge on chlorine increases** from ClO⁻ to ClO₄⁻ due to the increasing number of oxygen atoms attached to chlorine. This increased stability of the conjugate base leads to increased acidic strength of the corresponding acid ``.
  • Interhalogen Compounds:
    • Among interhalogen compounds, the maximum number of atoms are present in iodine fluoride ``.
    • Interhalogen compounds are more reactive than halogen compounds ``.
    • ClF₃ exists but FCl₃ does not because fluorine is more electronegative than chlorine, allowing it to form compounds like ClF₃ where chlorine is the central atom with higher oxidation states ``.
  • Tear Gas: CCl₃NO₂ is a tear gas ``.
• Noble Gases (Xenon)
  • Interactions: The only type of interactions between particles of noble gases are due to weak dispersion forces ``.
  • Ionisation Enthalpy: The ionisation enthalpy of molecular oxygen (O₂) is very close to that of xenon (Xe) ``. This similarity was key to Bartlett's work.
  • Bartlett's work: Bartlett had taken O₂⁺PtF₆⁻ as a base compound because both O₂ and Xe have almost the same ionisation enthalpy ``.
  • Reactivity of Xenon Fluorides: Xenon fluorides are reactive, contrary to the statement that they are not reactive ``.
  • Hydrolysis of XeF₆: The hydrolysis of XeF₆ is not a redox reaction . Its partial hydrolysis does not change the oxidation state of the central atom.
  • Uses: Helium (He) is used in modern diving apparatus . **Argon (Ar)** is used to provide an **inert atmosphere for filling electrical bulbs**.
• Other Important Concepts and Reactions
  • Qualitative Analysis (Copper Salt): When H₂S is passed through an aqueous solution of salt acidified with dilute HCl, a black precipitate is obtained. On boiling the precipitate with dilute HNO₃, it forms a solution of blue colour. Addition of excess of aqueous solution of ammonia to this solution gives a deep blue solution of [Cu(NH₃)₄]²⁺ ``.
  • Brown Ring Test for NO₃⁻ ion: A brown ring is formed due to the formation of [Fe(H₂O)₅(NO)]²⁺ ``. The reactions involved are:
    • NO₃⁻ + 3Fe²⁺ + 4H⁺ → NO + 3Fe³⁺ + 2H₂O ``
    • [Fe(H₂O)₆]²⁺ + NO → [Fe(H₂O)₅(NO)]²⁺ + H₂O (brown complex) ``
  • Oxidation State Change: In the reaction where a black compound of manganese reacts with a halogen acid to give greenish yellow gas (Cl₂), which then reacts with excess NH₃ to form an unstable trihalide, the oxidation state of nitrogen changes from –3 to +3 ``. (The option -3 to 0 is also plausible depending on the product, however, the source states -3 to +3).
  • Lead (II) Nitrate Decomposition: On heating lead (II) nitrate, it gives a **brown gas 'A' (NO₂) . Gas 'A' on cooling changes to a **colourless solid 'B' (N₂O₄). Solid 'B' on heating with NO changes to a **blue solid 'C' (N₂O₃) ``.
• Bond Angles: H₂O has a higher bond angle than H₂S . This is because oxygen is more electronegative than sulphur, causing the bond pair electrons of the O–H bond to be closer to oxygen, leading to **more bond-pair bond-pair repulsion** between the bond pairs of the two O–H bonds.

III. Molecular Structures and Shapes

• Cyclotrimetaphosphoric Acid Molecule: It has 3 double bonds and 9 single bonds ``.
• Bond Lengths:
  • All the three N—O bond lengths in HNO₃ are not equal ``.
  • All P—Cl bond lengths in PCl₅ molecule in gaseous state are not equal ``.
• Common Molecular Shapes and Hybridization:
  • NH₄⁺: Tetrahedral shape ``.
  • SiCl₄: Tetrahedral shape ``.
  • SF₄: Sea-saw shaped ``.
  • SO₄²⁻: Tetrahedral shape ``.
  • BrF₃: Bent T-shaped ``.
  • BrO₃⁻: Pyramidal shape ``.
  • XeF₆: sp³d³ – distorted octahedral hybridization and shape ``.
  • XeO₃: sp³ – pyramidal hybridization and shape ``.
  • XeOF₄: sp³d² – square pyramidal hybridization and shape ``.
  • XeF₄: sp³d² – square planar hybridization and shape . Its central atom is in sp³d² hybridization.

IV. Types of Oxides

• Pb₃O₄: Mixed oxide ``.
• N₂O: Neutral oxide ``.
• Mn₂O₇: Acidic oxide ``.
• Bi₂O₃: Basic oxide ``.

In essence, these p-block elements are like a diverse family of chemical compounds, each with unique personalities – some are strong reducers, some are powerful oxidizers, some are gases that change color with temperature, and others form complex structures, all governed by the subtle dance of electrons and atomic forces.

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