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JEE Mains Chemistry · Teaching notes

The p-Block Elements — JEE Mains Chemistry

The p-Block Elements has 185 past-year questions from 2021 to 2026, and 169 of them are multiple choice. It is a recall chapter: most questions are pairs of statements or lists of claims, and each claim is a trend or one of its exceptions, such as gallium being smaller than aluminium or chlorine beating fluorine in electron gain enthalpy. The few counting questions follow fixed rules: only a P–OH hydrogen ionises, and the lone pairs on an interhalogen's central atom follow from its formula. Check every statement on its own before looking at the options.

Every subtopic, worked example, formula and trap in one printable document — answers shown, ready to share.

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Subtopic notes

Formula & revision sheet

7 formulas · 15 reference tables · 58 gotchas across all subtopics — the exam-eve cheat-sheet

Group 13: Periodic Trends and the Inert Pair Effect

Reference tables (3)

Group 13 trends in radius, ionisation enthalpy and electronegativity5 rows
ElementAtomic radius (pm)M³⁺ radius (pm)First ionisation enthalpy (kJ/mol)Electronegativity
B85278012.0
Al14353.55771.5
Ga13562.05791.6
Smaller than Al and with a slightly higher ionisation enthalpy: poor shielding by 3d electrons.
In16780.05581.7
Tl17088.55891.8
NCERT values. Read each column on its own: the atomic radius dips at Ga, the M³⁺ radius does not, and the ionisation enthalpy is lowest at In, not Tl.
Group 13 melting points, boron's lattice and gallium's liquid range5 rows
ElementMelting point (K)Boiling point (K)Density (g/cm³)What to remember
B245339232.35Giant covalent B12\mathrm{B_{12}} network: very hard, highest melting point
Al93327402.70Light metal; made passive by concentrated HNO3\mathrm{HNO_3}, which coats it with oxide
Ga30326765.90Liquid from 303 K to 2676 K, the widest liquid range; used in high-temperature thermometers
The lowest melting point in the group, and still a liquid in boiling water.
In43023537.31Soft metal that melts above gallium
Tl576173011.85The densest member of the group
Melting points fall from boron to gallium and then rise a little: B > Al > Tl > In > Ga.
The inert pair effect in group 13: Tl⁺ is more stable than Tl³⁺4 rows
ElementMore stable oxidation stateE° for M³⁺ reduction (V)How M³⁺ behaves
Al+3 only−1.66-1.66 (Al3+/Al\mathrm{Al^{3+}/Al})Very stable; hard to reduce
Ga+3−0.56-0.56 (Ga3+/Ga\mathrm{Ga^{3+}/Ga})Stable; +1 appears only in salts such as GaAlCl4\mathrm{GaAlCl_4}
In+3−0.34-0.34 (In3+/In\mathrm{In^{3+}/In})Stable; In+\mathrm{In^{+}} is easily oxidised back to +3
Tl+1+1.26+1.26 (Tl3+\mathrm{Tl^{3+}} reduced to Tl+\mathrm{Tl^{+}})Strong oxidising agent
The positive potential is the inert pair effect in numbers: Tl³⁺ is eager to become Tl⁺.
A more positive reduction potential means the ion is more easily reduced, so a stronger oxidising agent.

Watch out for (8)

Boron and Aluminium Compounds

Formulas (1)

Reference tables (2)

Structure and preparation of diborane and borazine6 rows
FeatureDiborane, B₂H₆Borazine, B₃N₃H₆
ShapeNon-planar: the two BH2\mathrm{BH_2} ends lie in one plane, the two bridging H above and below itPlanar six-membered ring of alternating B and N
BondsFour terminal 2-centre-2-electron B–H bonds and two bridging 3-centre-2-electron B–H–B bondsOnly ordinary 2-centre-2-electron bonds, with π electrons delocalised round the ring
Banana bonds belong to diborane, never to borazine.
Hybridisation of boronAbout sp3sp^3sp2sp^2
Bond angles and lengthsTerminal H–B–H 122°, bridge H–B–H 97°; terminal B–H 119 pm, bridging B–H 134 pmAll six B–N bonds equal in length
With waterB2H6+6H2O→2B(OH)3+6H2\mathrm{B_2H_6 + 6H_2O \rightarrow 2B(OH)_3 + 6H_2}B3N3H6+9H2O→3B(OH)3+3NH3+3H2\mathrm{B_3N_3H_6 + 9H_2O \rightarrow 3B(OH)_3 + 3NH_3 + 3H_2}
Acid-base natureLewis acid; split by bases such as NMe3\mathrm{NMe_3}Polar B–N bonds make it more reactive than benzene
The terminal H–B–H angle is wider than the bridge angle, so the terminal bonds have more s character and less p character.
Boron and aluminium halides as Lewis acids: back-bonding and maximum covalency6 rows
SpeciesCovalency of the central atomShapeWhy
BF3\mathrm{BF_3}3Trigonal planarElectron deficient; back-bonding from F partly fills boron's empty p orbital
[BF4]−\mathrm{[BF_4]^{-}}4 (oxidation state still +3)TetrahedralFluoride donates a pair into boron's empty orbital
BF63−\mathrm{BF_6^{3-}}Would need 6Does not existBoron has no d orbitals, so four bonds is its limit
The reason NCERT gives is the missing d orbitals.
[AlF6]3−\mathrm{[AlF_6]^{3-}}6OctahedralAluminium uses its 3d orbitals
[Al(H2O)6]3+\mathrm{[Al(H_2O)_6]^{3+}}6Octahedral, sp3d2sp^3d^2Formed when aluminium chloride dissolves in acidified water
Al2Cl6\mathrm{Al_2Cl_6}4Two tetrahedra sharing an edge of two bridging ClEach aluminium completes its octet through a chlorine lone pair
Covalency counts the bonds round the atom; it is not the oxidation state. Boron is +3 in both BF₃ and [BF₄]⁻.

Watch out for (7)

Group 14: Carbon, Silicon, Tin and Lead

Reference tables (3)

Group 14 trends, carbon's allotropes and silicones5 rows
ElementCovalent radius (pm)First ionisation enthalpy (kJ/mol)ElectronegativityWhat sets it apart
C7710862.5Catenation and pπp\pi–pπp\pi bonds; maximum covalency 4; allotropes
Si1187861.8Uses d orbitals: [SiF6]2−\mathrm{[SiF_6]^{2-}} exists; SiO2\mathrm{SiO_2} is acidic; forms silicones
Ge1227611.8GeO2\mathrm{GeO_2} acidic; [GeCl6]2−\mathrm{[GeCl_6]^{2-}} exists
Sn1407081.8+4 more stable than +2; oxides amphoteric
Pb1467151.9+2 more stable than +4; oxides amphoteric
Lead's ionisation enthalpy is a little HIGHER than tin's: poor shielding by 4f and 5d electrons.
The ionisation enthalpy falls from C to Sn, then rises slightly at Pb.
Inert pair effect in tin and lead: which ions oxidise and which reduce6 rows
IonPreferred state of the elementBehaves asEvidence
Sn2+\mathrm{Sn^{2+}}+4Reducing agentE∘(Sn4+/Sn2+)=+0.15 V\mathrm{E^\circ(Sn^{4+}/Sn^{2+}) = +0.15\ V}: Sn2+\mathrm{Sn^{2+}} is easily oxidised
Sn4+\mathrm{Sn^{4+}}+4Stable; a very weak oxidant at mostSame small potential, +0.15 V+0.15\ V
Pb2+\mathrm{Pb^{2+}}+2StableThe 6s pair stays out of bonding
Pb4+\mathrm{Pb^{4+}}+2Strong oxidising agentE∘(Pb4+/Pb2+)=+1.67 V\mathrm{E^\circ(Pb^{4+}/Pb^{2+}) = +1.67\ V}, the most positive here
The strongest oxidant among these p-block ions.
Tl3+\mathrm{Tl^{3+}}+1Strong oxidising agentTl3+\mathrm{Tl^{3+}} reduced to Tl+\mathrm{Tl^{+}}: +1.26 V+1.26\ V
Tl+\mathrm{Tl^{+}}+1StableThe 6s pair stays out of bonding
The more positive the reduction potential, the stronger the oxidising agent.
Tests for the lead ion in salt analysis5 rows
Reagent added to Pb²⁺ProductColourWhat happens next
Dilute HClPbCl2\mathrm{PbCl_2}WhiteDissolves on heating the water
H2S\mathrm{H_2S}PbS\mathrm{PbS}BlackDissolves in hot dilute HNO3\mathrm{HNO_3} to give Pb(NO3)2\mathrm{Pb(NO_3)_2}
K2CrO4\mathrm{K_2CrO_4}PbCrO4\mathrm{PbCrO_4}YellowDissolves in NaOH as Na2[Pb(OH)4]\mathrm{Na_2[Pb(OH)_4]}
Charge 2−, four OH groups: coordination number 4.
KIPbI2\mathrm{PbI_2}YellowDissolves in hot water and returns as golden spangles on cooling
Dilute H2SO4\mathrm{H_2SO_4}PbSO4\mathrm{PbSO_4}WhiteDissolves in ammonium acetate solution
Chloride, sulphate and nitrate of lead are the white or colourless ones; chromate and iodide are yellow; sulphide is black.

Watch out for (7)

Group 15: Periodic Trends and Hydrides

Reference tables (2)

Group 15 trends and the anomalous behaviour of nitrogen5 rows
ElementCovalent radius (pm)First ionisation enthalpy (kJ/mol)ElectronegativityCharacter
N7014023.0Non-metal, diatomic gas N2\mathrm{N_2}
No d orbitals: maximum covalency 4.
P11010122.1Non-metal, P4\mathrm{P_4} molecules
As1219472.0Metalloid
Sb1418341.9Metalloid
Bi1487031.9Metal, the only one in the group
The biggest steps are between N and P in every column; below arsenic the changes are small.
Hydrides of group 15 from NH₃ to BiH₃4 rows
HydrideH–E–H angle (°)Boiling point (K)E–H bond enthalpy (kJ/mol)Character
NH3\mathrm{NH_3}107.8238.5389Most stable and most basic; weakest reducing agent; hydrogen bonded
PH3\mathrm{PH_3}93.6185.5322Lowest boiling point in the group: no hydrogen bonding and a small molar mass
The lowest boiling point is PH₃, not NH₃.
AsH3\mathrm{AsH_3}91.8210.6297Less basic and more reducing than PH3\mathrm{PH_3}
SbH3\mathrm{SbH_3}91.3254.6255Highest boiling point of the four: the largest dispersion forces
BiH₃, not listed because it is too unstable to measure well, continues every trend: least stable, least basic, strongest reducing agent.

Watch out for (6)

Nitrogen and Its Compounds

Formulas (1)

Reference tables (1)

Oxides of nitrogen: oxidation states, structures and nature6 rows
OxideOxidation state of NStructureNature
N2O\mathrm{N_2O}+1Linear N≡N–O; one N–N bondNeutral; colourless gas
NO\mathrm{NO}+2N=O with one unpaired electronNeutral; colourless gas
N2O3\mathrm{N_2O_3}+3O=N–NO₂; one N–N bondAcidic; blue solid
NO2\mathrm{NO_2}+4Bent, odd electron on N; one N=O and one N–OAcidic; brown gas
The odd-electron oxide that dimerises to N₂O₄.
N2O4\mathrm{N_2O_4}+4O₂N–NO₂; one N–N bond, no bridging OAcidic; colourless
N2O5\mathrm{N_2O_5}+5O₂N–O–NO₂; one N–O–N bridge, no N–N bondAcidic; colourless solid, the anhydride of HNO3\mathrm{HNO_3}
Only the two lowest oxides are neutral; only N₂O₅ bridges its nitrogens through oxygen.

Watch out for (5)

Phosphorus and Its Oxoacids

Formulas (2)

  • Allotropes of phosphorus and reactions of white phosphorus and its chlorides · Key reactions of phosphorus
    P4+3NaOH+3H2O→PH3+3NaH2PO2P4+8SOCl2→4PCl3+4SO2+2S2Cl2PCl3+3H2O→H3PO3+3HClPCl5+4H2O→H3PO4+5HCl\mathrm{P_4 + 3NaOH + 3H_2O \rightarrow PH_3 + 3NaH_2PO_2} \qquad \mathrm{P_4 + 8SOCl_2 \rightarrow 4PCl_3 + 4SO_2 + 2S_2Cl_2} \qquad \mathrm{PCl_3 + 3H_2O \rightarrow H_3PO_3 + 3HCl} \qquad \mathrm{PCl_5 + 4H_2O \rightarrow H_3PO_4 + 5HCl}
  • Basicity and reducing power of phosphorus oxoacids · Counting rule for phosphorus oxoacids
    basicity=number of P−OH groupsnon-ionisable H=number of P−H bondsP−H present⇒reducing\text{basicity} = \text{number of } \mathrm{P{-}OH} \text{ groups} \qquad \text{non-ionisable H} = \text{number of } \mathrm{P{-}H} \text{ bonds} \qquad \mathrm{P{-}H} \text{ present} \Rightarrow \text{reducing}

Reference tables (1)

Oxoacids of phosphorus: formulas, oxidation states and bonds8 rows
AcidFormulaOxidation state of PBonds in the structure
Hypophosphorous (phosphinic)H3PO2\mathrm{H_3PO_2}+1Two P–H, one P–OH, one P=O
Orthophosphorous (phosphonic)H3PO3\mathrm{H_3PO_3}+3One P–H, two P–OH, one P=O
PyrophosphorousH4P2O5\mathrm{H_4P_2O_5}+3Two P–H, two P–OH, two P=O, one P–O–P
HypophosphoricH4P2O6\mathrm{H_4P_2O_6}+4One P–P, four P–OH, two P=O
Hypophosphoric (+4, P–P bond) is not hypophosphorous (+1, two P–H).
OrthophosphoricH3PO4\mathrm{H_3PO_4}+5Three P–OH, one P=O
PyrophosphoricH4P2O7\mathrm{H_4P_2O_7}+5Four P–OH, two P=O, one P–O–P
Cyclotrimetaphosphoric(HPO3)3\mathrm{(HPO_3)_3}+5A ring with three P–O–P, three P–OH, three P=O
Phosphorus(V) oxideP4O10\mathrm{P_4O_{10}}+5Six P–O–P bridges and four P=O (the anhydride, not an acid)
Pyrophosphorous is +3 with P–H bonds; pyrophosphoric is +5 with none.

Watch out for (7)

Group 16: Oxygen and Sulphur

Formulas (1)

  • Redox reactions of sulphur compounds and the tests for sulphide and sulphite · Sulphur redox reactions
    Cr2O72−+3SO2+2H+→2Cr3++3SO42−+H2OS8+12OH−→4S2−+2S2O32−+6H2O2S2O32−+I2→S4O62−+2I−\mathrm{Cr_2O_7^{2-} + 3SO_2 + 2H^{+} \rightarrow 2Cr^{3+} + 3SO_4^{2-} + H_2O} \qquad \mathrm{S_8 + 12OH^{-} \rightarrow 4S^{2-} + 2S_2O_3^{2-} + 6H_2O} \qquad \mathrm{2S_2O_3^{2-} + I_2 \rightarrow S_4O_6^{2-} + 2I^{-}}

Reference tables (2)

Group 16 trends: oxygen's anomalies, hydrides and oxides4 rows
HydrideMelting point (K)H–E bond enthalpy (kJ/mol)H–E–H angle (°)Acid strength (Ka)
H2O\mathrm{H_2O}2734631041.8×10−161.8 \times 10^{-16}
Hydrogen bonding makes water melt highest, though it is the lightest.
H2S\mathrm{H_2S}188347921.3×10−71.3 \times 10^{-7}
H2Se\mathrm{H_2Se}208276911.3×10−41.3 \times 10^{-4}
H2Te\mathrm{H_2Te}222238902.3×10−32.3 \times 10^{-3}
The bond enthalpy falls down the group, so acid strength and reducing power rise: H₂Te is the strongest acid and strongest reducing agent of the four.
Oxoacids of sulphur: structures, S=O bonds and oxidation states7 rows
AcidFormulaOxidation state of SS=O bondsLink between units
SulphurousH2SO3\mathrm{H_2SO_3}+41One unit; a lone pair on S
SulphuricH2SO4\mathrm{H_2SO_4}+62One unit, two S–OH
ThiosulphuricH2S2O3\mathrm{H_2S_2O_3}Average +2; the two S differ1A terminal S doubly bonded to the central S, in place of one O
DithionicH2S2O6\mathrm{H_2S_2O_6}+5, both S alike4A direct S–S bond
Pyrosulphuric (oleum)H2S2O7\mathrm{H_2S_2O_7}+64One S–O–S bridge
Peroxodisulphuric (Marshall's)H2S2O8\mathrm{H_2S_2O_8}+64One O–O peroxo bridge
Still +6: the two peroxo oxygens are −1 each.
PolythionicH2SxO6\mathrm{H_2S_xO_6}Ends +5, chain 04A chain of S atoms between two SO3H\mathrm{SO_3H} groups
Each S=O bond carries one π bond, so counting S=O counts the π bonds of every acid here except thiosulphuric, which also has an S=S.

Watch out for (9)

Groups 17 and 18: Halogens and Noble Gases

Formulas (2)

Reference tables (1)

Halogen properties: bond enthalpy, electron gain enthalpy and hydrogen halides4 rows
Halogen (hydride)X–X bond enthalpy (kJ/mol)Electron gain enthalpy (kJ/mol)HX boiling point (K)HX melting point (K)
F (HF)158.8−333-333293190
Weak F–F bond and a less negative electron gain enthalpy than Cl: both from fluorine's small size.
Cl (HCl)242.6−349-349189159
Br (HBr)192.8−325-325206185
I (HI)151.1−296-296238222
Chlorine leads in both bond enthalpy and electron gain enthalpy. HF boils highest; HI melts highest.

Watch out for (9)

PYQ weightage by concept

22 concepts · 185 PYQs — where the marks actually sit, so you know what to drill first

Group 13: Periodic Trends and the Inert Pair Effect25 PYQs · 14%
ConceptPYQsShare
Group 13 trends in radius, ionisation enthalpy and electronegativity95%
The inert pair effect in group 13: Tl⁺ is more stable than Tl³⁺95%
Group 13 melting points, boron's lattice and gallium's liquid range74%
Boron and Aluminium Compounds21 PYQs · 11%
ConceptPYQsShare
Borax, the borax bead test and boric acid95%
Structure and preparation of diborane and borazine74%
Boron and aluminium halides as Lewis acids: back-bonding and maximum covalency53%
Group 14: Carbon, Silicon, Tin and Lead16 PYQs · 9%
ConceptPYQsShare
Group 14 trends, carbon's allotropes and silicones63%
Tests for the lead ion in salt analysis63%
Inert pair effect in tin and lead: which ions oxidise and which reduce42%
Group 15: Periodic Trends and Hydrides21 PYQs · 11%
ConceptPYQsShare
Group 15 trends and the anomalous behaviour of nitrogen126%
Hydrides of group 15 from NH₃ to BiH₃95%
Nitrogen and Its Compounds18 PYQs · 10%
ConceptPYQsShare
Oxides of nitrogen: oxidation states, structures and nature95%
Preparing dinitrogen and nitric acid, and the nitrogen tests95%
Phosphorus and Its Oxoacids19 PYQs · 10%
ConceptPYQsShare
Basicity and reducing power of phosphorus oxoacids74%
Allotropes of phosphorus and reactions of white phosphorus and its chlorides63%
Oxoacids of phosphorus: formulas, oxidation states and bonds63%
Group 16: Oxygen and Sulphur30 PYQs · 16%
ConceptPYQsShare
Group 16 trends: oxygen's anomalies, hydrides and oxides137%
Redox reactions of sulphur compounds and the tests for sulphide and sulphite95%
Oxoacids of sulphur: structures, S=O bonds and oxidation states84%
Groups 17 and 18: Halogens and Noble Gases35 PYQs · 19%
ConceptPYQsShare
Oxidising power and disproportionation of the halogens137%
Interhalogen shapes, halogen oxoacids and xenon fluorides137%
Halogen properties: bond enthalpy, electron gain enthalpy and hydrogen halides95%

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