From the air we breathe to the chips in your phone — Groups 15–18 of the periodic table. With an interactive p-block element explorer, structural diagrams of P₄, S₈ and XeF₄, and 20+ real-life examples.
Before diving in, here is the whole chapter on one page. Four groups, each with its own trends, allotropes, and key compounds.
The p-block covers Groups 13–18; this chapter focuses on Groups 15–18. All have the outer configuration ns²np¹⁻⁶. Key trends: atomic radius increases down a group, ionisation energy and electronegativity decrease. Group 15: N, P, As, Sb, Bi. Group 16: O, S, Se, Te, Po. Group 17: halogens. Group 18: noble gases. Each group has characteristic allotropes and important compounds, and each shows interesting anomalies due to small size and absence of d-orbitals.
p-Block elements: Elements in which the last electron enters a p-orbital of the outermost shell. The general outer configuration is ns² np¹⁻⁶. They occupy Groups 13–18 of the modern periodic table.
Groups covered in this chapter:
| Property | Down the group | Reason |
|---|---|---|
| Atomic radius | Increases | New shells added |
| Ionisation energy | Decreases | Outer electron farther from nucleus, more shielding |
| Electronegativity | Decreases | Same as above |
| Metallic character | Increases | Easier to lose electrons |
| Density | Increases (mostly) | More mass packed in a similar volume |
| Stability of highest oxidation state | Decreases | Inert-pair effect — heavier elements prefer lower oxidation states by 2 |
Elements of the second period (N, O, F) differ from the rest of their group because:
Inert-pair effect: The reluctance of the outermost s-electrons to take part in bonding as we go down a p-block group. It makes the lower oxidation state (group number minus 2) more stable for heavier elements.
p-Block = outer ns²np¹⁻⁶. Group 15: ns²np³. Group 16: ns²np⁴. Group 17: ns²np⁵. Group 18: ns²np⁶. Down a group: radius ↑, IE ↓, EN ↓, metallic character ↑. Second-period elements (N, O, F) differ because of small size, high EN and no d-orbitals. Inert-pair effect makes group number minus 2 stable for heavier elements.
Elements: Nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), bismuth (Bi).
Electronic configuration: ns² np³ — 5 valence electrons. They need 3 more to complete the octet, so their common oxidation states are −3, +3, +5.
Occurrence:
| Property | Trend | Note |
|---|---|---|
| Atomic radius | Increases N → Bi | N < P < As < Sb < Bi |
| Ionisation energy | Decreases | N highest, Bi lowest |
| Electronegativity | Decreases | N (3.0) > P (2.1) > ... > Bi (1.9) |
| Metallic character | Increases | N and P non-metals, As metalloid, Sb and Bi metals |
| Melting point | Decreases | N₂ is a gas; Bi is a solid |
| Common oxidation states | −3, +3, +5 | Inert-pair effect stabilises +3 down the group |
Properties: Colourless, odourless, non-toxic gas. Very unreactive because the two N atoms are held by a very strong triple bond (N≡N) with high bond enthalpy (~941 kJ/mol).
Preparation (in the laboratory):
Uses: Inert atmosphere in the manufacture of chemicals and in packaging; liquid nitrogen as a coolant; fertiliser manufacture (Haber process).
1. White phosphorus (P₄):
2. Red phosphorus:
3. Black phosphorus:
Group 15: N, P, As, Sb, Bi. Configuration ns²np³; oxidation states −3, +3, +5. N₂ is inert due to N≡N. P has white (reactive, P₄, stored under water), red (polymeric, stable, safety matches) and black (most stable, layered) allotropes. Metallic character increases down the group.
Structure: Pyramidal. N is sp³ hybridised with one lone pair. Bond angle ~107°.
Preparation (Haber process):
N₂ + 3H₂ ⇌ 2NH₃ (Fe catalyst, 723 K, 200 atm)
Laboratory preparation: 2NH₄Cl + Ca(OH)₂ → 2NH₃ + CaCl₂ + 2H₂O
Properties: Colourless, pungent gas, very soluble in water (forms NH₄OH). Basic — reacts with acids to give ammonium salts. Forms a white precipitate with Nessler's reagent. Acts as a Lewis base (donates the lone pair).
Uses: Fertilisers (urea, ammonium nitrate), nitric acid manufacture, refrigeration, cleaning agents, laboratory reagent.
Preparation (Ostwald process):
Properties: Colourless liquid (commercial acid is yellow because of dissolved NO₂), strong monobasic acid, powerful oxidising agent.
Reactions:
Uses: Fertilisers (ammonium nitrate), explosives (TNT, nitrocellulose), etching metals, rocket propellants, oxidising agent in the laboratory.
PCl₃: Pyramidal (sp³, one lone pair). Prepared by passing dry Cl₂ over white P. Hydrolyses to H₃PO₃ + HCl. Used to make PCl₅, SOCl₂ and as a chlorinating agent.
PCl₅: Trigonal bipyramidal in the gas and liquid state (sp³d hybridisation), but ionic [PCl₄]⁺[PCl₆]⁻ in the solid state. Prepared by reacting PCl₃ with excess Cl₂. Hydrolyses to POCl₃ + HCl, and then to H₃PO₄. Used as a chlorinating agent.
| Acid | Formula | Oxidation state of P | Basicity |
|---|---|---|---|
| Hypophosphorous | H₃PO₂ | +1 | Monobasic |
| Phosphorous | H₃PO₃ | +3 | Dibasic |
| Orthophosphoric | H₃PO₄ | +5 | Tribasic |
| Pyrophosphoric | H₄P₂O₇ | +5 | Tetrabasic |
| Metaphosphoric | HPO₃ | +5 | Monobasic |
Basicity: Number of replaceable H atoms (attached to O). H₃PO₂ has 1 (only one OH); H₃PO₃ has 2 (two OH); H₃PO₄ has 3 (three OH).
HNO₂ (nitrous acid): Weak monobasic acid, unstable, disproportionates. N oxidation state +3.
HNO₃ (nitric acid): Strong monobasic acid, powerful oxidiser. N oxidation state +5.
Don't confuse basicity with number of H atoms. H₃PO₃ has 3 hydrogen atoms but is dibasic because only 2 of the H atoms are attached to oxygen (the third is directly attached to P and is not acidic). H₃PO₂ has 3 H atoms but is monobasic.
NH₃ — Haber process, pyramidal, basic. HNO₃ — Ostwald process, strong oxidiser, makes fertilisers and explosives. PCl₃ — pyramidal; PCl₅ — trigonal bipyramidal (gas) but ionic in the solid state. Oxoacids of P: H₃PO₂ (monobasic, P +1), H₃PO₃ (dibasic, P +3), H₃PO₄ (tribasic, P +5). Basicity counts only OH hydrogens.
Elements: Oxygen (O), sulphur (S), selenium (Se), tellurium (Te), polonium (Po).
Electronic configuration: ns² np⁴ — 6 valence electrons. They need 2 more to complete the octet, so their common oxidation states are −2, +2, +4, +6.
Occurrence: O₂ makes up ~21% of air and 46% of the Earth's crust. S occurs as elemental sulphur and as metal sulphides (ZnS, PbS, CuFeS₂) and sulphates (CaSO₄·2H₂O).
| Property | Trend | Note |
|---|---|---|
| Atomic radius | Increases | O < S < Se < Te < Po |
| Ionisation energy | Decreases | O highest, Po lowest |
| Electronegativity | Decreases | O (3.5) > S (2.5) > Se > Te |
| Metallic character | Increases | O and S non-metals, Se and Te metalloids, Po a metal |
| Common oxidation states | −2, +2, +4, +6 | Inert-pair effect stabilises +4 down the group |
Structure: Bent molecule (bond angle ~117°). Resonance between two structures gives O–O bond order 1.5.
Preparation: Silent electrical discharge through dry oxygen:
3O₂ → 2O₃ (silent electric discharge, 10% conversion)
Properties:
Ozone hole: CFCs (Freons) release Cl atoms in the stratosphere, which catalytically destroy ozone:
The Cl atom is regenerated — one Cl can destroy thousands of ozone molecules. The Montreal Protocol banned CFCs.
Uses: Water sterilisation, air purification, bleaching oils and textiles.
Rhombic sulphur (α-sulphur): Yellow, stable below 369 K. S₈ crown-shaped molecules.
Monoclinic sulphur (β-sulphur): Stable above 369 K, needle-shaped crystals.
Plastic sulphur: Formed by pouring molten sulphur into cold water. Non-crystalline, elastic, used in rubber vulcanisation.
All solid allotropes are based on the crown-shaped S₈ molecule.
Group 16: O, S, Se, Te, Po. Configuration ns²np⁴; oxidation states −2, +2, +4, +6. Ozone (O₃) is bent, a powerful oxidiser, and protects Earth from UV — destroyed by CFCs. Sulphur allotropes are all based on the crown-shaped S₈ molecule; rhombic and monoclinic are the main crystalline forms.
Preparation: Burning sulphur in air, or roasting metal sulphides:
Structure: Bent molecule (bond angle ~119°). S is sp² hybridised with one lone pair.
Properties: Colourless, pungent gas. Acidic oxide — dissolves in water to give sulphurous acid H₂SO₃. Reducing agent (converts Fe³⁺ to Fe²⁺, decolourises KMnO₄). Bleaches by reduction (temporary bleaching).
Uses: Manufacture of H₂SO₄, bleaching wool and silk, preservative, disinfectant, refrigerant.
Preparation (Contact process):
Note: SO₃ is absorbed in concentrated H₂SO₄ (not water) to avoid a violent reaction and acid mist.
Properties: Colourless, oily liquid, highly corrosive. Strong diprotic acid. Powerful oxidising and dehydrating agent.
Dehydrating agent: Removes water from sugar (turns it black — "charring"): C₁₂H₂₂O₁₁ → 12C + 11H₂O. Also removes water from ethanol, oxalic acid, etc.
Oxidising agent: Cu + 2H₂SO₄ (conc.) → CuSO₄ + SO₂ + 2H₂O
Uses: Fertilisers (superphosphate, ammonium sulphate), lead-acid batteries, petroleum refining, dyes, detergents, explosives, metallurgy, laboratory reagent.
| Acid | Formula | Oxidation state of S |
|---|---|---|
| Sulphurous acid | H₂SO₃ | +4 |
| Sulphuric acid | H₂SO₄ | +6 |
| Pyrosulphuric acid (oleum) | H₂S₂O₇ | +6 |
| Thiosulphuric acid | H₂S₂O₃ | +6 and −2 (one S each) |
| Peroxodisulphuric acid | H₂S₂O₈ | +6 |
SO₂ — bent, acidic, reducing, bleaches by reduction. H₂SO₄ — made by the Contact process (V₂O₅ catalyst); strong diprotic acid, powerful oxidiser and dehydrating agent. SO₃ is absorbed in concentrated H₂SO₄, not water. H₂SO₄ is used in fertilisers, batteries, petroleum refining and dyes.
Elements: Fluorine (F), chlorine (Cl), bromine (Br), iodine (I), astatine (At).
Electronic configuration: ns² np⁵ — 7 valence electrons. They need just 1 more to complete the octet, so they are highly reactive and form −1 ions easily (halides).
Occurrence: F as fluorite CaF₂, cryolite Na₃AlF₆, fluoroapatite; Cl as NaCl (sea water), carnallite; Br and I in sea water and brine.
State at room temperature:
| Property | Trend | Note |
|---|---|---|
| Atomic radius | Increases | F < Cl < Br < I |
| Ionisation energy | Decreases | F highest |
| Electronegativity | Decreases | F (4.0) the most electronegative element |
| Electron affinity | Decreases (Cl > F) | F is anomalously low due to small size, electron repulsion |
| Oxidising power | Decreases | F₂ > Cl₂ > Br₂ > I₂ |
| Bond dissociation enthalpy | Cl₂ > Br₂ > F₂ > I₂ | F₂ anomalously low due to lone-pair repulsion |
HF: Weak acid (H-bonding, high bond enthalpy). Attacks glass: SiO₂ + 4HF → SiF₄ + 2H₂O. Used for etching glass.
HCl: Strong acid. Prepared by treating NaCl with concentrated H₂SO₄. Colourless, pungent gas, fumes in moist air.
HBr and HI: Stronger acids than HCl, but reducing — HI is the strongest reducing agent among the HX.
Acid strength: HF < HCl < HBr < HI (opposite to bond enthalpy — weaker bond = stronger acid).
| Acid | Formula | Halogen oxidation state |
|---|---|---|
| Hypochlorous | HClO | +1 |
| Chlorous | HClO₂ | +3 |
| Chloric | HClO₃ | +5 |
| Perchloric | HClO₄ | +7 |
Acid strength: HClO < HClO₂ < HClO₃ < HClO₄ — more oxygen atoms increase acidity by stabilising the conjugate base.
Interhalogens: Compounds formed between two different halogens. Types: AX (ClF, BrF, ICl, IBr), AX₃ (ClF₃, BrF₃, ICl₃), AX₅ (ClF₅, BrF₅, IF₅), AX₇ (IF₇).
Structure: AX — linear. AX₃ — T-shaped. AX₅ — square pyramidal. AX₇ — pentagonal bipyramidal. (Increasing number of lone pairs changes the geometry.)
Reactivity: Greater than the halogens themselves; F forms compounds with all other halogens.
Use: ClF₃ and BrF₃ are used as fluorinating agents; ICl is used in organic synthesis.
Preparation: Cl₂ + Ca(OH)₂ → Ca(OCl)Cl + H₂O
Bleaching action: Ca(OCl)Cl + H₂O → Ca(OH)₂ + Cl₂ (or HOCl). The chlorine/hypochlorous acid bleaches by oxidation.
Uses: Water disinfection, bleaching cotton and paper, as a disinfectant.
Group 17: F, Cl, Br, I, At. Configuration ns²np⁵; mostly −1 oxidation state, but +1 to +7 in oxoacids. Oxidising power: F₂ > Cl₂ > Br₂ > I₂. Acid strength: HF < HCl < HBr < HI. Oxoacids HClO < HClO₂ < HClO₃ < HClO₄. Interhalogens: AX, AX₃, AX₅, AX₇. Bleaching powder = Ca(OCl)Cl.
Elements: Helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), radon (Rn).
Electronic configuration: ns² np⁶ (a closed octet) — except He, which is 1s². This closed shell makes them chemically inert and gives them very high ionisation energies.
Occurrence: In the atmosphere (Ar ~0.93%, Ne ~0.0018%, He ~0.0005%, Kr ~0.0001%, Xe ~0.000009%). He is also found in natural gas.
State at room temperature: All are monatomic gases.
| Property | Trend | Note |
|---|---|---|
| Atomic radius | Increases | He < Ne < Ar < Kr < Xe < Rn |
| Ionisation energy | Decreases | He highest, Rn lowest |
| Boiling point | Increases | He lowest, Rn highest (dispersion forces increase) |
| Solubility in water | Increases | He least, Xe most |
| Chemical reactivity | Increases | He, Ne, Ar inert; Kr and Xe form compounds |
Xenon fluorides:
Preparation: Xe + excess F₂ (1:5 ratio) at 873 K and 7 bar gives XeF₆; 1:2 ratio at 673 K gives XeF₄; 1:1 ratio at 673 K and 1 bar gives XeF₂.
Xenon oxides:
Xenon oxyfluorides: XeOF₄, XeO₂F₂, XeO₃F₂.
For decades, chemists believed the noble gases could not form any compounds at all. The first noble-gas compound, XePtF₆, was made by Neil Bartlett in 1962. Since then, hundreds of xenon compounds (and some krypton compounds) have been synthesised — but helium, neon and argon remain completely inert under ordinary conditions.
Group 18: He, Ne, Ar, Kr, Xe, Rn. Closed-shell ns²np⁶ configuration makes them inert (He 1s²). Reactivity increases down the group — Kr and Xe form compounds, He/Ne/Ar do not. XeF₂ linear, XeF₄ square planar, XeF₆ distorted octahedral. Uses: He in balloons and MRI, Ne in signs, Ar in welding, Xe in headlights.
| Element / Compound | Applications |
|---|---|
| N₂ | Inert atmosphere, food packaging, liquid N₂ coolant |
| NH₃ | Fertilisers, nitric acid, refrigeration |
| HNO₃ | Fertilisers, explosives, metal etching |
| H₃PO₄ | Fertilisers, food acid, rust remover |
| O₂ | Medical O₂, steel making, welding |
| O₃ | Water sterilisation, air purification |
| H₂SO₄ | Fertilisers, batteries, refining, dyes |
| SO₂ | H₂SO₄ manufacture, bleaching, preservative |
| Cl₂ | Water disinfection, PVC manufacture, bleach |
| HCl | Pickling steel, laboratory reagent, PVC |
| HF | Glass etching, semiconductor industry |
| I₂ | Antiseptic, iodised salt, photography |
| He | Balloons, MRI cooling, breathing mixtures |
| Ne | Advertising signs (red glow) |
| Ar | Welding shield gas, light bulbs |
| Xe | HID headlights, ion thrusters, anaesthesia |
1. Acid rain: SO₂ and NOₓ from burning fossil fuels dissolve in rain to give H₂SO₄ and HNO₃. Harms buildings, soil, forests, lakes.
2. Ozone depletion: CFCs release Cl atoms that catalytically destroy stratospheric ozone. Banned by the Montreal Protocol.
3. Photochemical smog: NOₓ and hydrocarbons react in sunlight to give ozone and PAN — harms lungs and plants.
4. Eutrophication: Excess nitrate and phosphate fertilisers wash into lakes and rivers, causing algal blooms that deplete oxygen.
5. Greenhouse gases: CO₂, and trace amounts of N₂O, contribute to global warming.
p-Block compounds are central to modern life — fertilisers, batteries, disinfectants, lighting, medical gases. Their misuse causes acid rain, ozone depletion, smog and eutrophication. Remedies: flue-gas desulphurisation, catalytic converters, CFC phase-out and controlled fertiliser use.
"N P As Sb Bi" — Never Put Any Salt Before Bread.
"O S Se Te Po" — Oh See Some Tea Pots.
"F Cl Br I At" — Funny Clowns Bring Ice Cream And Tea.
"He Ne Ar Kr Xe Rn" — He Never Argues, Krypton X-rays Remember Nothing.
"−3, +3, +5" — "Three three five" (group number − 10, group number, group number − 10).
"−2, +2, +4, +6" — "Two two four six" — start at group number − 10, then +2 at a time.
"−1, +1, +3, +5, +7" — "One one three five seven" — odd numbers, starting from −1.
"More O, stronger acid" — HClO < HClO₂ < HClO₃ < HClO₄.
"Longer bond = stronger acid" — HF < HCl < HBr < HI (opposite of what many expect).
"Count the OH groups, not all H atoms" — H₃PO₂ = 1 OH (monobasic), H₃PO₃ = 2 OH (dibasic), H₃PO₄ = 3 OH (tribasic).
"2-3-4-5-6-7 → linear-square-T-BP-oct" — XeF₂ linear (2 bonds, sp³d), XeF₄ square planar (4 bonds, sp³d²), XeF₆ distorted octahedral (6 bonds, sp³d³). Remember: lone pairs force the shape.
Use "Never Put Any Salt Before Bread" for Group 15, "Oh See Some Tea Pots" for Group 16, "Funny Clowns Bring Ice Cream" for Group 17, and "He Never Argues" for Group 18. Remember: more oxygen = stronger oxoacid; longer HX bond = stronger acid.
Q1. Why is nitrogen much less reactive than phosphorus?
N₂ has a very strong N≡N triple bond (bond enthalpy ~941 kJ/mol). It is a small, diatomic molecule with a very high bond dissociation energy, so it is kinetically inert.
Phosphorus exists as P₄ with P–P single bonds. Single bonds are much weaker than a triple bond (P–P bond enthalpy ~201 kJ/mol). Also, the P–P–P bond angle in the P₄ tetrahedron is only 60°, creating angle strain. Both factors make P₄ much more reactive.
Q2. Why is HF a weaker acid than HCl, even though F is more electronegative than Cl?
HF has a very high H–F bond enthalpy (~568 kJ/mol) due to the small size of F. Breaking this bond is difficult, so HF does not release H⁺ easily.
Fluoride ion (F⁻) is small and strongly H-bonded to water, which stabilises the undissociated HF form. HCl, HBr and HI do not show such strong H-bonding. In HCl, the H–Cl bond is much weaker (~431 kJ/mol), so it dissociates more easily.
Because bond enthalpy decreases down the group and the conjugate base stability increases (I⁻ best stabilised by delocalisation in water), the order is HF < HCl < HBr < HI. The dominant factor is the bond enthalpy, which decreases as the halogen atom gets larger.
Q3. Compare the acid strength of the oxoacids of chlorine.
HClO < HClO₂ < HClO₃ < HClO₄ (hypochlorous < chlorous < chloric < perchloric). Perchloric acid is the strongest.
As the number of oxygen atoms increases, the conjugate base is stabilised by greater delocalisation of the negative charge over more oxygen atoms. A more stable conjugate base means a stronger acid.
The oxidation state of Cl increases from +1 (HClO) to +7 (HClO₄). Higher positive charge on Cl pulls electrons away from the O–H bond, making it more polar and easier to ionise. Together these make HClO₄ the strongest.
Q4. Describe the preparation of ammonia by the Haber process. What are its uses?
N₂ + 3H₂ ⇌ 2NH₃ ΔH = −92 kJ/mol (exothermic).
(i) Catalyst: Finely divided iron with
molybdenum as a promoter.
(ii)
Temperature: ~723 K (compromise — too low
gives a slow rate; too high reduces yield due to the
exothermic nature).
(iii) Pressure: ~200
atm (high pressure favours the side with fewer gas moles —
that is NH₃).
(iv) The ammonia is removed as it forms,
shifting the equilibrium to the right (Le Chatelier).
NH₃ is pyramidal (sp³, one lone pair, bond angle ~107°). It is a colourless, pungent gas, very soluble in water (forming NH₄OH), basic, and a Lewis base.
(i) Fertilisers — urea, ammonium nitrate, ammonium
sulphate.
(ii) Manufacture of nitric acid (Ostwald
process).
(iii) Refrigeration.
(iv) Cleaning
agents.
(v) Laboratory reagent.
Q5. Explain the structure and bonding in PCl₅. Why is it ionic in the solid state but molecular in the gas phase?
In the gas and liquid phases, PCl₅ exists as discrete molecules. P is sp³d hybridised, giving a trigonal bipyramidal geometry (three equatorial P–Cl bonds, two axial P–Cl bonds). The axial bonds are slightly longer than the equatorial bonds due to greater repulsion.
In the solid state, PCl₅ exists as the ionic salt [PCl₄]⁺[PCl₆]⁻. The cation [PCl₄]⁺ is tetrahedral (P is sp³). The anion [PCl₆]⁻ is octahedral (P is sp³d²). The ionic form is favoured because lattice energy stabilises it.
In the solid state, the electrostatic lattice energy of the ionic pair makes the ionic arrangement more stable. In the gas phase, the ion pairs are not held together by the lattice — the covalent molecule is more stable because of the lower energy of the discrete molecular state. The interconversion between the molecular and ionic forms depends on the state.
Q6. Discuss the properties and uses of noble gases.
All noble gases are monatomic, colourless, odourless and tasteless. They have a closed-shell ns²np⁶ configuration (He is 1s²), giving very high ionisation energies and very low boiling points. They are essentially chemically inert, though Xe and Kr can form compounds with F and O under forcing conditions.
Down the group: atomic radius increases, ionisation energy decreases, boiling point increases (stronger dispersion forces), solubility in water increases. Reactivity also increases down the group.
Balloons and airships (light, inert — no fire risk), cryogenic coolant for MRI magnets, breathing mixtures for deep-sea divers, and leak detection.
Neon: advertising signs (red-orange glow), high-voltage indicators. Argon: inert shield gas in welding, filling incandescent bulbs, and inert atmosphere in metallurgy. Xenon: car headlights (HID), ion thrusters in spacecraft, anaesthesia, and flash lamps. Krypton: photographic flashes and energy-efficient lamps. Radon: radioactive — hazardous in some basements.
Selected exercise questions from the LEGACY NCERT Chapter 7 (The p-Block Elements). Removed from the current syllabus but useful for competitive exams.
In pentahalides (PX₅), the central atom (P, As, Sb) is in the +5 oxidation state and forms five covalent bonds. A higher oxidation state means a greater polarising power of the central atom, which pulls electron density towards itself. This reduces the ionic character of the bonds and makes them more covalent.
In trihalides (PX₃), the central atom is in the +3 oxidation state. The polarising power is lower, so the bonds have more ionic character.
Additionally, in pentahalides the central atom has a larger share of positive charge, which decreases the size of the cation and increases its polarising power — favouring covalent bonding (Fajans' rules).
N₂: The two nitrogen atoms are held together by a very strong N≡N triple bond (bond enthalpy ~941 kJ/mol). It is a small, diatomic molecule with no angle strain. The bond is also thermodynamically very stable because of the small size of N and effective overlap of p-orbitals.
P₄: Phosphorus exists as P₄ with weak P–P single bonds (bond enthalpy ~201 kJ/mol). The tetrahedral P₄ structure has a very small bond angle (60°), creating considerable angle strain. Both the weak bonds and the strain make P₄ chemically much more reactive.
Therefore, N₂ is a relatively inert gas under ordinary conditions, while P₄ (white phosphorus) is so reactive that it catches fire in air.
Both NH₃ and H₂O have sp³ hybridised central atoms, with bond angles less than the ideal tetrahedral angle of 109.5° because of lone-pair repulsion.
NH₃: N has one lone pair. The bond angle is ~107°.
H₂O: O has two lone pairs. Two lone pairs repel each other and the bonding pairs more strongly than one lone pair does, pushing the O–H bonds closer together. The bond angle is ~104.5°.
So the order is H₂O (104.5°) < NH₃ (107°) < CH₄ (109.5°). This is the standard explanation based on VSEPR theory.
Laboratory preparation: 2NH₄Cl + Ca(OH)₂ → 2NH₃ + CaCl₂ + 2H₂O. The ammonia gas is dried by passing through quicklime (CaO) — not through concentrated H₂SO₄, which would react with it.
(a) Excess Cl₂: 8NH₃ + 3Cl₂ → 6NH₄Cl + N₂. Ammonia is oxidised to nitrogen, and the excess chlorine gives white fumes of NH₄Cl. (With excess NH₃ and limited Cl₂, the product is N₂H₄ or NH₂Cl instead.)
(b) Nessler's reagent (K₂HgI₄ in KOH): Gives a brown precipitate (the "iodide of Millon's base") — a sensitive test for ammonia: 2K₂HgI₄ + 3KOH + NH₃ → [Hg₂NI]·H₂O (brown ppt) + 7KI + H₂O.
(a) H₃PO₃ (phosphorous acid): Two OH groups and one P–H bond: (HO)₂P(=O)H. Only the two OH hydrogens are acidic, so it is dibasic.
(b) H₃PO₂ (hypophosphorous acid): One OH group and two P–H bonds: (HO)P(=O)H₂. Only one OH hydrogen is acidic, so it is monobasic.
(c) H₃PO₄ (orthophosphoric acid): Three OH groups: (HO)₃P(=O). All three hydrogens are acidic, so it is tribasic.
Rule: basicity = number of OH groups, not the total number of H atoms.
SO₂: Bent molecule. The S atom is sp² hybridised with one lone pair, giving a bond angle of ~119°. Two resonance structures give the S=O bonds partial double-bond character. The molecule is polar.
SO₃: Trigonal planar in the gas phase. The S atom is sp² hybridised, forming three S=O bonds at 120° to each other. In the solid state it forms polymeric chains (as (SO₃)ₙ) or trimeric (SO₃)₃ rings. The molecule is non-polar in the gas phase.
Why "king of chemicals": Sulphuric acid is the most widely produced industrial chemical in the world. Its uses range from fertilisers (superphosphate, ammonium sulphate) to lead-acid batteries, petroleum refining, dyes, detergents and explosives. Any country's level of industrialisation is often judged by its sulphuric acid consumption.
Contact process:
Step 1 — Production of SO₂: S + O₂ → SO₂, or by roasting sulphide ores (4FeS₂ + 11O₂ → 2Fe₂O₃ + 8SO₂).
Step 2 — Oxidation of SO₂: 2SO₂ + O₂ ⇌ 2SO₃. Conditions: V₂O₅ catalyst, 720 K, 2 atm. The reaction is exothermic and reversible; 720 K is a compromise between reaction rate and yield.
Step 3 — Absorption: SO₃ + H₂SO₄ → H₂S₂O₇ (oleum). SO₃ is absorbed in concentrated H₂SO₄, not in water, to avoid a violent reaction and acid mist.
Step 4 — Dilution: H₂S₂O₇ + H₂O → 2H₂SO₄.
(a) Oxidising power: F₂ > Cl₂ > Br₂ > I₂. Fluorine is the strongest oxidising agent among the halogens because of its very low bond dissociation enthalpy and high electron affinity (its small size makes electron gain very favourable).
(b) Acid strength of HX: HF < HCl < HBr < HI. The acid strength is governed by the H–X bond enthalpy, which decreases down the group. So HI is the strongest acid even though iodine is the least electronegative halogen.
(c) Stability of oxoacids: For a given oxidation state, the oxoacids of lighter halogens are generally more stable. For example, HClO₄ is more stable than HBrO₄ or HIO₄. Within one halogen, higher oxidation state oxoacids are stronger and more stable (HClO₄ > HClO₃ > HClO₂ > HClO).
Monatomic: Noble gases have a closed-shell ns²np⁶ configuration (He is 1s²). This configuration is very stable — there is no tendency to gain or lose electrons, so they do not form covalent bonds with each other. Each atom is therefore independent, and they exist as single atoms.
Low boiling points: Because they are monatomic and have no permanent dipole moment, the only force between noble gas atoms is the weak London dispersion force. Down the group, these forces increase (larger, more polarisable atoms) so boiling point increases from He to Rn, but overall they remain very low compared to other elements of similar mass. This is why all noble gases are gases at room temperature.
XeF₂: Xe is sp³d hybridised with 3 lone pairs. The 3 lone pairs occupy the equatorial positions of a trigonal bipyramid, leaving the 2 F atoms in axial positions. The result is a linear molecule (bond angle 180°).
XeF₄: Xe is sp³d² hybridised with 2 lone pairs. The 2 lone pairs occupy trans positions of an octahedron, leaving the 4 F atoms in the equatorial plane. The result is a square planar molecule.
XeF₆: Xe is sp³d³ hybridised with 1 lone pair. The 1 lone pair distorts the octahedron, pushing the 6 F atoms out of the ideal positions. The result is a distorted octahedral (or capped octahedral) structure. The lone pair can also participate in rapid internal rearrangements.
In all three, the number of lone pairs determines the shape (VSEPR).
| Group 15 valence e⁻ | 5 (ns² np³) | Oxidation states −3, +3, +5 |
| Group 16 valence e⁻ | 6 (ns² np⁴) | Oxidation states −2, +2, +4, +6 |
| Group 17 valence e⁻ | 7 (ns² np⁵) | Oxidation states −1, +1, +3, +5, +7 |
| Group 18 valence e⁻ | 8 (ns² np⁶) | Closed shell — inert |
| N≡N bond enthalpy | ~941 kJ/mol | Very strong — N₂ inert |
| P–P bond enthalpy | ~201 kJ/mol | Weak — P₄ reactive |
| H–F bond enthalpy | ~568 kJ/mol | Strong — HF weak acid |
| H–I bond enthalpy | ~299 kJ/mol | Weak — HI strong acid |
| Acid strength of HX | HF < HCl < HBr < HI | Bond enthalpy decreases |
| Oxidising power of halogens | F₂ > Cl₂ > Br₂ > I₂ | F₂ strongest oxidiser |
| Acid strength of Cl oxoacids | HClO < HClO₂ < HClO₃ < HClO₄ | More O = stronger |
| XeF₂ geometry | Linear (sp³d) | 3 lone pairs equatorial |
| XeF₄ geometry | Square planar (sp³d²) | 2 lone pairs trans |
| XeF₆ geometry | Distorted octahedral (sp³d³) | 1 lone pair |
| ns² np³ | Group 15 outer configuration |
| ns² np⁴ | Group 16 outer configuration |
| ns² np⁵ | Group 17 outer configuration |
| ns² np⁶ | Group 18 outer configuration |
| HX | Hydrogen halide (HF, HCl, HBr, HI) |
| PX₃ / PX₅ | Phosphorus trihalide / pentahalide |
| H₂SO₄ | Sulphuric acid |
| H₂S₂O₇ | Pyrosulphuric acid (oleum) |
| Ca(OCl)Cl | Bleaching powder |
| H₃PO₂ / H₃PO₃ / H₃PO₄ | Hypophosphorous / phosphorous / orthophosphoric acid |
(A) Both true, R explains A. (B) Both true, R does NOT explain A. (C) A true, R false. (D) A false, R true.
All key equations and facts in one place — perfect for last-minute revision before competitive exams.