Chapter 7 — The p-Block Elements — covers Groups 15, 16, 17 and 18 of the periodic table. This is a fact-heavy chapter with 8-10 marks in Boards, making it one of the highest-weightage chapters. Focus on the preparation and properties of key compounds: NH₃, HNO₃, H₂SO₄, ozone, and interhalogen compounds.
Key Concepts
Group 15 — Nitrogen Family (N, P, As, Sb, Bi)
General electronic configuration: ns²np³
Trends Down the Group
- Metallic character increases: N, P (non-metals) → As, Sb (metalloids) → Bi (metal)
- Oxidation states: −3 to +5; stability of +5 state decreases and +3 increases down the group (inert pair effect)
- Bi shows +3 more stable than +5 (inert pair effect)
Ammonia (NH₃)
Conditions: 450-500°C, 200 atm, Fe catalyst + Mo promoter
Lab preparation: NH₄Cl + Ca(OH)₂ → CaCl₂ + 2H₂O + 2NH₃↑
Properties:
- Colourless gas with pungent smell, highly soluble in water (forms NH₄OH)
- Basic nature: donates lone pair → acts as Lewis base and Brønsted base
- Forms complexes: Cu²⁺ + 4NH₃ → [Cu(NH₃)₄]²⁺ (deep blue — test for Cu²⁺)
Nitric Acid (HNO₃)
4NH₃ + 5O₂ →(Pt/Rh, 500°C) 4NO + 6H₂O
2NO + O₂ → 2NO₂
3NO₂ + H₂O → 2HNO₃ + NO
Properties: Strong oxidising acid. Dissolves all metals except Au and Pt. Makes metals passive (concentrated HNO₃ makes Fe, Cr, Al passive by forming oxide layer).
Au + 3HCl + HNO₃ → HAuCl₄ + NO + 2H₂O
Oxides of Nitrogen
| Oxide | Name | Nature | Oxidation State of N |
|---|---|---|---|
| N₂O | Nitrous oxide (laughing gas) | Neutral | +1 |
| NO | Nitric oxide | Neutral | +2 |
| N₂O₃ | Dinitrogen trioxide | Acidic | +3 |
| NO₂ | Nitrogen dioxide | Acidic | +4 |
| N₂O₅ | Dinitrogen pentoxide | Acidic | +5 |
Allotropes of Phosphorus
| White P | Red P | Black P |
|---|---|---|
| P₄ tetrahedral | Polymeric chains | Layered structure |
| Waxy, poisonous, glows in dark | Non-poisonous, doesn’t glow | Most stable allotrope |
| Soluble in CS₂ | Insoluble in CS₂ | Insoluble in CS₂ |
| Catches fire at 35°C | Ignites at 260°C | Most unreactive |
Group 16 — Oxygen Family (O, S, Se, Te, Po)
General electronic configuration: ns²np⁴
Ozone (O₃)
Structure: Angular/V-shaped, O−O bond length = 128 pm (between single and double bond), resonance hybrid
Properties: Powerful oxidising agent, decomposes to O₂ on heating, decolourises KMnO₄, liberates I₂ from KI (used as test).
Sulphuric Acid (H₂SO₄) — “King of Chemicals”
S + O₂ → SO₂
2SO₂ + O₂ →(V₂O₅, 450°C) 2SO₃
SO₃ + H₂SO₄ → H₂S₂O₇ (oleum)
H₂S₂O₇ + H₂O → 2H₂SO₄
Properties:
- Dehydrating agent: Removes water from compounds (chars sugar, formic acid)
- Oxidising agent: Hot conc. H₂SO₄ oxidises metals, non-metals, and compounds
- Dibasic acid: H₂SO₄ → H⁺ + HSO₄⁻ → 2H⁺ + SO₄²⁻
Oxoacids of Sulphur
| Acid | Formula | Key Feature |
|---|---|---|
| Sulphurous acid | H₂SO₃ | Reducing agent |
| Sulphuric acid | H₂SO₄ | Strong acid, dehydrating |
| Thiosulphuric acid | H₂S₂O₃ | S replaces one O in H₂SO₄ |
| Peroxodisulphuric acid | H₂S₂O₈ | Contains S−O−O−S (peroxide linkage) |
Group 17 — Halogens (F, Cl, Br, I)
General electronic configuration: ns²np⁵
Trends
- Most electronegative elements; electronegativity decreases down group
- Oxidising power: F₂ > Cl₂ > Br₂ > I₂
- All show −1 oxidation state; Cl, Br, I also show +1, +3, +5, +7 (F never shows positive states)
- Bond dissociation energy: Cl₂ > Br₂ > F₂ > I₂ (F₂ unexpectedly low due to small size → lone pair repulsion)
Hydrogen Halides (HX)
| Property | HF | HCl | HBr | HI |
|---|---|---|---|---|
| Acidic strength | Weakest | ↓ | ↓ | Strongest |
| Thermal stability | Most stable | ↓ | ↓ | Least stable |
| Reducing power | Least | ↓ | ↓ | Most |
| Boiling point | High (H-bonding) | 189 K | 206 K | 238 K |
Interhalogen Compounds
| Type | Shape | Examples |
|---|---|---|
| XX’ | Linear | ClF, BrF, ICl, IBr |
| XX’₃ | T-shaped | ClF₃, BrF₃, ICl₃ |
| XX’₅ | Square pyramidal | BrF₅, IF₅ |
| XX’₇ | Pentagonal bipyramidal | IF₇ |
Group 18 — Noble Gases (He, Ne, Ar, Kr, Xe, Rn)
Fully filled orbitals → very stable, low reactivity.
Xenon Compounds
| Compound | Hybridisation | Shape | Oxidation State |
|---|---|---|---|
| XeF₂ | sp³d | Linear | +2 |
| XeF₄ | sp³d² | Square planar | +4 |
| XeF₆ | sp³d³ | Distorted octahedral | +6 |
| XeO₃ | sp³ | Pyramidal | +6 |
| XeOF₂ | sp³d | T-shaped | +4 |
Important Definitions
| Term | Definition |
|---|---|
| Inert Pair Effect | Reluctance of s-electrons to participate in bonding in heavier elements |
| Allotropy | Existence of an element in two or more forms with different physical properties |
| Interhalogen Compound | Compound formed between two different halogen atoms |
| Ozone | Triatomic allotrope of oxygen (O₃), powerful oxidising agent |
| Aqua Regia | Mixture of 3:1 conc. HCl and HNO₃ that dissolves gold and platinum |
Solved Examples — NCERT Based
Example 1: Structures of Oxoacids
Q: Why does H₃PO₃ (phosphorous acid) act as a diprotic acid despite having 3 hydrogen atoms?
Solution: In H₃PO₃, only 2 hydrogen atoms are attached to oxygen (O−H bonds) and can be ionised. The third hydrogen is directly bonded to phosphorus (P−H bond), which is not ionisable. Hence it is diprotic (dibasic), not triprotic.
Example 2: Oxidation States
Q: Why does BiH₃ not exist while NH₃ is very stable?
Solution: Due to the inert pair effect, the 6s² electrons of Bi are reluctant to participate in bonding. The Bi−H bond is very weak (large size difference), making BiH₃ extremely unstable. NH₃ is stable because of strong N−H bonds (small size, good overlap).
Example 3: Halogen Reactivity
Q: Why is F₂ the strongest oxidising agent among halogens?
Solution: Three factors favour F₂:
1. Low F−F bond dissociation energy (easy to break)
2. High electronegativity (strong tendency to gain electrons)
3. Very high hydration enthalpy of F⁻ (small size → strong hydration)
All three factors make the overall ΔG very negative for reduction, making F₂ the strongest oxidising agent.
Example 4: Xenon Compound Geometry
Q: Predict the shape and hybridisation of XeF₄.
Solution: Xe has 8 valence electrons. In XeF₄, 4 bond pairs + 2 lone pairs = 6 electron pairs → sp³d² hybridisation. The 2 lone pairs occupy opposite positions (axial) to minimise repulsion → shape is square planar.
Important Questions for Board Exams
1 Mark Questions
- Why is N₂ less reactive at room temperature?
- What is the basicity of H₃PO₄?
- Why does fluorine not show positive oxidation states?
- What is the shape of XeF₂?
- Why is HF a weak acid despite F being most electronegative?
2 Mark Questions
- Draw the structures of (a) H₂SO₄ (b) HNO₃.
- Explain why Bi³⁺ is more stable than Bi⁵⁺.
- Why is SO₂ an air pollutant?
- What is the structure of ozone? Is it a resonance hybrid?
3 Mark Questions
- Explain the Contact process for manufacturing H₂SO₄ with equations.
- Discuss the preparation and properties of ammonia.
- What are interhalogen compounds? Give their types, examples, and shapes.
- How does HNO₃ react with (a) copper (b) zinc (c) iron?
5 Mark Questions
- Discuss the chemistry of Group 16 elements. Explain the preparation of H₂SO₄ by Contact process. What are its properties as (a) dehydrating agent (b) oxidising agent?
- Describe the preparation and properties of XeF₂, XeF₄ and XeF₆. Write their structures.
Quick Revision Points
- Group 15: ns²np³; NH₃ (Haber), HNO₃ (Ostwald); inert pair effect → Bi³⁺ stable
- Group 16: ns²np⁴; O₃ (angular), H₂SO₄ (Contact process — “King of Chemicals”)
- Group 17: ns²np⁵; strongest oxidisers; F₂ never +ve; interhalogen compounds
- Group 18: noble gases; Xe forms compounds (XeF₂ linear, XeF₄ sq. planar, XeF₆ distorted oct.)
- Aqua Regia = 3HCl : 1HNO₃ (dissolves Au, Pt)
- H₃PO₃ is dibasic (not tribasic) — one P−H bond
- HF is a weak acid due to high H−F bond strength despite F’s electronegativity
- F₂ bond energy < Cl₂ bond energy (lone pair repulsion in small F₂)
Chapter Navigation
Previous: General Principles and Processes of Isolation of Elements Class 12 Notes
Next: The d and f Block Elements Class 12 Notes
Related Chapters in Class 12 Chemistry
- The d and f Block Elements Class 12 Notes
- General Principles and Processes of Isolation of Elements Class 12 Notes
- Coordination Compounds Class 12 Notes
Practice What You Learned
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