JEE Mains Chemistry · Teaching notes

Chemical Bonding and Molecular Structure — JEE Mains Chemistry

Chemical Bonding and Molecular Structure has 176 past-year questions from 2021 to 2026, and 54 of them ask for a number rather than an option. Most of those numbers are counts: how many species in a list are linear, polar or paramagnetic, how many lone pairs sit on a central atom, how many σ and π bonds a chain holds. Each count rests on a short routine done the same way every time, so a careful count of electrons and electron pairs is worth more here than any formula. The rest of the chapter is orders and exceptions worth knowing by heart, such as NF₃ against NH₃ or O₂⁺ against O₂⁻.

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

All 176 past questions JEE Mains Chemistry strategy JEE Mains home

Subtopic notes

Formula & revision sheet

8 formulas · 12 reference tables · 42 gotchas across all subtopics — the exam-eve cheat-sheet

Lewis Structures, Formal Charge and the Octet Rule

Formulas (1)

Counting lone pairs and formal charge · Lone pairs and formal charge $\text{lone pairs} = \frac{N_{\text{valence}} - 2\,n_{\text{bonds}}}{2} \qquad \text{FC} = V - L - \tfrac{1}{2}S$

Reference tables (2)

The octet rule and its three exceptions5 rows

Type	Electrons on the central atom	Examples
Obeys the octet rule	8	$\mathrm{CH_4}$ , $\mathrm{CO_2}$ , $\mathrm{CCl_4}$ , $\mathrm{NH_3}$ , $\mathrm{SiF_4}$ , $\mathrm{H_2S}$
Incomplete octet	4 for Be, 6 for B and Al	$\mathrm{BeF_2}$ , $\mathrm{BeH_2}$ (4); $\mathrm{BF_3}$ , $\mathrm{BCl_3}$ , $\mathrm{AlCl_3}$ (6)
Electron deficient	Bridging B–H–B bonds hold 2 electrons over 3 atoms	$\mathrm{B_2H_6}$ ; $\mathrm{BCl_3}$ is also called electron deficient
Odd-electron species	An odd total: NO 11, NO₂ 17, ClO₂ 19	$\mathrm{NO}$ , $\mathrm{NO_2}$ , $\mathrm{ClO_2}$ These are also the paramagnetic oxides: an odd electron cannot pair.
Expanded octet	10 or 12 (14 in IF₇)	$\mathrm{PCl_5}$ , $\mathrm{SF_4}$ (10); $\mathrm{SF_6}$ , $\mathrm{H_2SO_4}$ , $\mathrm{SO_3}$ (12); $\mathrm{IF_7}$ (14)

Only period 3 and heavier atoms can expand the octet; N, O and F never do.

Lewis acids and Lewis bases6 rows

Species	Role	Reason
$\mathrm{BF_3}$ , $\mathrm{BCl_3}$	Lewis acid	B has 6 electrons and an empty p orbital; $sp^2$ , trigonal planar
$\mathrm{AlCl_3}$	Lewis acid	Al has 6 electrons; it dimerises to $\mathrm{Al_2Cl_6}$ to fill the gap
$\mathrm{BI_3}$	Strongest boron halide acid	Back-bonding from large I into B is weakest
$\mathrm{NH_3}$ , $\mathrm{NF_3}$	Lewis base	One lone pair on N; $sp^3$ , pyramidal
$\mathrm{SF_4}$ , $\mathrm{ClF_3}$	Lewis base	One lone pair on S; two on Cl
$\mathrm{PCl_5}$	Not a Lewis base	All five P electrons are in bonds; no lone pair PCl₅ can accept a pair (forming PCl₆⁻), but it cannot donate one.

An empty orbital makes an acid; a lone pair on the central atom makes a base.

Watch out for (6)

Count every atom, not just the centre
→ Counting lone pairs and formal charge
The charge changes the electron count
→ Counting lone pairs and formal charge
Sulphur acids and oxides are expanded
→ The octet rule and its three exceptions
Odd electrons mean an exception, even with a small atom
→ The octet rule and its three exceptions
Electronegativity does not rank the boron halides
→ Lewis acids and Lewis bases
Amphoteric water is a Brønsted idea
→ Lewis acids and Lewis bases

Ionic Bonding, Lattice Enthalpy and Fajans' Rules

Formulas (1)

The Born-Haber cycle and lattice enthalpy · Born-Haber cycle $\Delta_f H = \Delta_{sub}H + \Delta_i H + \tfrac{1}{2}\Delta_{diss}H + \Delta_{eg}H + \Delta_{lattice}H$

Reference tables (1)

Fajans' rules and covalent character5 rows

Rule	Order of covalent character	Why
Smaller cation	$\mathrm{LiCl > NaCl > KCl > CsCl}$	$\mathrm{Li^+}$ is the smallest and most polarising
Higher cation charge	$\mathrm{AlCl_3 > MgCl_2 > NaCl}$ ; $\mathrm{SnCl_4 > SnCl_2}$	More charge on a smaller ion
Larger anion	$\mathrm{CaI_2 > CaBr_2 > CaCl_2 > CaF_2}$ ; $\mathrm{KI > KF}$	$\mathrm{I^-}$ has the largest, softest cloud
18-electron cation	$\mathrm{CuCl > NaCl}$ ; $\mathrm{AgCl > KCl}$	d electrons shield the nuclear charge poorly
Electronegativity difference	Ionic character: $\mathrm{N_2 < ClF_3 < SO_2 < K_2O < LiF}$	$\Delta\chi$ is 0 for $\mathrm{N_2}$ , about 0.8 for Cl–F, 0.9 for S–O

The same polarisation that adds covalent character lowers the melting point and the solubility in water.

Watch out for (4)

Half the bond enthalpy, not all of it
→ The Born-Haber cycle and lattice enthalpy
Check which way the lattice step runs
→ The Born-Haber cycle and lattice enthalpy
A bigger cation means LESS covalent
→ Fajans' rules and covalent character
Rank electron gain by magnitude
→ Fajans' rules and covalent character

Bond Length, Bond Angle and Resonance

Formulas (1)

Resonance and fractional bond order · Bond order in a resonance hybrid $\text{bond order} = \frac{\text{total bonds to the equivalent atoms}}{\text{number of equivalent atoms}}$

Reference tables (2)

Bond length and what sets it10 rows

Bond	Typical length (pm)	Note
C–H	109	Shortest here: hydrogen is tiny
C≡C	120	Triple bond
C=C	134	Double bond
C–C	154	Single bond
C≡N	116	Shorter than C=O despite N being larger than C
C=O	122	Carbonyl
C–O	143	Alcohols and ethers
O=O	121	In O₂
O–O	148	In H₂O₂
P–Cl in PCl₅	219 axial, 204 equatorial	Axial bonds are the longer, weaker pair Calling the axial bonds of PCl₅ stronger is a standard wrong statement.

For the same pair of atoms: triple shorter than double, double shorter than single.

Bond angles and lone-pair repulsion8 rows

Species	Pairs on the centre	Bond angle
$\mathrm{BF_3}$	3 bond, 0 lone	120°
$\mathrm{SO_2}$	2 bond (plus π), 1 lone	about 119°
$\mathrm{CH_4}$	4 bond, 0 lone	109.5°
$\mathrm{NH_3}$	3 bond, 1 lone	107°
$\mathrm{H_2O}$	2 bond, 2 lone	104.5°
$\mathrm{NF_3}$	3 bond, 1 lone	102°
$\mathrm{PF_3}$	3 bond, 1 lone	about 98°
$\mathrm{ClF_3}$	3 bond, 2 lone	about 87.5° (axial F–Cl–equatorial F) The two lone pairs bend the axial F atoms back below 90°.

More lone pairs on the centre, or more electronegative outer atoms, means a smaller angle.

Watch out for (6)

The hybrid does not flip between forms
→ Resonance and fractional bond order
Resonance, not repulsion, sets ozone's bond length
→ Resonance and fractional bond order
Bond length is not set by bond order alone
→ Bond length and what sets it
See-saw and trigonal bipyramid give unequal bonds
→ Bond length and what sets it
Both SO₂ and H₂O are bent, at very different angles
→ Bond angles and lone-pair repulsion
Fluorine closes the angle; chlorine opens it
→ Bond angles and lone-pair repulsion

VSEPR Theory and Lone Pairs

Formulas (1)

Counting lone pairs on the central atom · Steric number and lone pairs $SN = \tfrac{1}{2}(V + M - c + a) \qquad \text{lone pairs} = SN - (\text{atoms bonded to the centre})$

Reference tables (1)

Where lone pairs sit: equatorial and trans7 rows

Species	Bond pairs, lone pairs	Lone pairs sit	Shape
$\mathrm{SF_4}$ , $\mathrm{SeF_4}$	4, 1	Equatorial	See-saw
$\mathrm{ClF_3}$ , $\mathrm{BrF_3}$	3, 2	Both equatorial	T-shaped (bent T), about 87.5°
$\mathrm{XeF_2}$ , $\mathrm{I_3^-}$ , $\mathrm{ICl_2^-}$	2, 3	All three equatorial	Linear, 180°
$\mathrm{XeO_2F_2}$	4, 1	Equatorial, with the two O	See-saw, F atoms axial
$\mathrm{BrF_5}$ , $\mathrm{IF_5}$	5, 1	Any one octahedral site	Square pyramidal
$\mathrm{XeF_4}$ , $\mathrm{BrF_4^-}$	4, 2	Trans, opposite each other	Square planar, 90°
$\mathrm{BrF_2^+}$	2, 2	Two corners of a tetrahedron	Bent

Five pairs: lone pairs equatorial. Six pairs: two lone pairs trans.

Watch out for (4)

A double bond to oxygen uses two electrons
→ Counting lone pairs on the central atom
Know what the question calls a bond pair
→ Counting lone pairs on the central atom
Lone pairs are never axial in a trigonal bipyramid
→ Where lone pairs sit: equatorial and trans
Three lone pairs make a straight molecule
→ Where lone pairs sit: equatorial and trans

Shapes of Molecules and Ions

Reference tables (1)

Shapes from the AXE formula15 rows

Type	Pairs (X + E)	Shape	Examples
AX₂	2	Linear	$\mathrm{BeCl_2}$ , $\mathrm{CO_2}$ , $\mathrm{NO_2^+}$ , $\mathrm{N_3^-}$ , $\mathrm{HC{\equiv}C^-}$
AX₃	3	Trigonal planar	$\mathrm{BF_3}$ , $\mathrm{SO_3}$ , $\mathrm{NO_3^-}$ , $\mathrm{CO_3^{2-}}$
AX₂E	3	Bent	$\mathrm{SO_2}$ , $\mathrm{O_3}$ , $\mathrm{NO_2^-}$
AX₄	4	Tetrahedral	$\mathrm{CH_4}$ , $\mathrm{NH_4^+}$ , $\mathrm{SO_4^{2-}}$ , $\mathrm{SO_2Cl_2}$
AX₃E	4	Trigonal pyramidal	$\mathrm{NH_3}$ , $\mathrm{H_3O^+}$ , $\mathrm{SO_3^{2-}}$ , $\mathrm{ClO_3^-}$ , $\mathrm{BrO_3^-}$ , $\mathrm{XeO_3}$
AX₂E₂	4	Bent	$\mathrm{H_2O}$ , $\mathrm{OF_2}$ , $\mathrm{ClO_2^-}$ , $\mathrm{BrF_2^+}$
AX₅	5	Trigonal bipyramidal	$\mathrm{PCl_5}$ , $\mathrm{PF_5}$ , $\mathrm{Fe(CO)_5}$
AX₄E	5	See-saw	$\mathrm{SF_4}$ , $\mathrm{SeF_4}$ , $\mathrm{IF_4^+}$ , $\mathrm{XeO_2F_2}$
AX₃E₂	5	T-shaped	$\mathrm{ClF_3}$ , $\mathrm{BrF_3}$ , $\mathrm{IF_3}$
AX₂E₃	5	Linear	$\mathrm{XeF_2}$ , $\mathrm{I_3^-}$ , $\mathrm{IBr_2^-}$
AX₆	6	Octahedral	$\mathrm{SF_6}$ , $\mathrm{[CrF_6]^{3-}}$
AX₅E	6	Square pyramidal	$\mathrm{BrF_5}$ , $\mathrm{IF_5}$ , $\mathrm{XeOF_4}$
AX₄E₂	6	Square planar	$\mathrm{XeF_4}$ , $\mathrm{ICl_4^-}$ , $\mathrm{BrF_4^-}$
AX₇	7	Pentagonal bipyramidal	$\mathrm{IF_7}$
AX₆E	7	Distorted octahedral	$\mathrm{XeF_6}$

Same total of pairs, same arrangement; the lone pairs decide the name of the shape.

Watch out for (4)

Name the shape from the atoms, not the pairs
→ Shapes from the AXE formula
Same formula type, different shape
→ Shapes from the AXE formula
Charge changes the shape
→ Counting species of one shape in a list
Two nickel complexes, two shapes
→ Counting species of one shape in a list

Hybridisation and Sigma and Pi Bonds

Formulas (2)

Counting σ and π bonds · Counting bonds in an open chain $n_\sigma = N_{\text{atoms}} - 1 \ (+1 \text{ per ring}) \qquad n_\pi = n_{\text{double}} + 2\,n_{\text{triple}}$
Hybridisation from the steric number · Steric number to hybridisation $SN = n_\sigma + n_{\text{lp}}: \ 2 \to sp,\ 3 \to sp^2,\ 4 \to sp^3,\ 5 \to sp^3d,\ 6 \to sp^3d^2,\ 7 \to sp^3d^3$

Reference tables (1)

Hybridisation, orientation and complexes8 rows

Hybridisation	Orientation	Main-group examples	Complex examples
$sp$	Linear, 180°	$\mathrm{BeCl_2}$ , $\mathrm{CO_2}$ , $\mathrm{NO_2^+}$	$\mathrm{[Ag(NH_3)_2]^+}$
$sp^2$	Trigonal planar, 120°	$\mathrm{BF_3}$ , $\mathrm{SO_2}$ , $\mathrm{NO_2^-}$	Rare in complexes
$sp^3$	Tetrahedral, 109.5°	$\mathrm{CH_4}$ , $\mathrm{NH_4^+}$ , $\mathrm{XeO_3}$	$\mathrm{Ni(CO)_4}$ , $\mathrm{[NiCl_4]^{2-}}$
$dsp^2$	Square planar, 90°	None	$\mathrm{[PtCl_4]^{2-}}$ , $\mathrm{[Ni(CN)_4]^{2-}}$
$sp^3d$	Trigonal bipyramidal	$\mathrm{PCl_5}$ , $\mathrm{SF_4}$ , $\mathrm{ClF_3}$ , $\mathrm{XeF_2}$	$\mathrm{Fe(CO)_5}$ is often written $dsp^3$
$sp^3d^2$	Octahedral, 90°	$\mathrm{SF_6}$ , $\mathrm{BrF_5}$ , $\mathrm{XeF_4}$	$\mathrm{[CoF_6]^{3-}}$ (outer orbital)
$d^2sp^3$	Octahedral, 90°	None	$\mathrm{[Co(NH_3)_6]^{3+}}$ , $\mathrm{[Fe(CN)_6]^{3-}}$ (inner orbital) [Co(NH₃)₆]³⁺ is d²sp³, not sp³d²: a stated match of it with SF₆ is false.
$sp^3d^3$	Pentagonal bipyramidal	$\mathrm{IF_7}$ , $\mathrm{XeF_6}$ (distorted)	None

The hybridisation fixes the arrangement of pairs; the shape then depends on how many are lone pairs.

Watch out for (6)

Every C–H bond is a σ bond
→ Counting σ and π bonds
Give σ and π in the order asked
→ Counting σ and π bonds
π bonds do not add hybrid orbitals
→ Hybridisation from the steric number
Five atoms around the centre can still be sp³d²
→ Hybridisation from the steric number
sp³d with a lone pair gives unequal bonds
→ Hybridisation, orientation and complexes
Square planar means dsp², not sp³
→ Hybridisation, orientation and complexes

Molecular Orbital Theory

Formulas (1)

Bond order from the MO diagram · Bond order $\text{bond order} = \tfrac{1}{2}(N_b - N_a)$

Reference tables (2)

Combining atomic orbitals (LCAO)8 rows

Pair of orbitals (axis z)	Symmetry of each	Do they combine?
1s and 1s	σ and σ	Yes: σ1s and σ*1s
$2p_z$ and $2p_z$	σ and σ	Yes: σ2p and σ*2p (head-on)
$2p_x$ and $2p_x$	π and π	Yes: π2p and π*2p (sideways)
2s and $2p_z$	σ and σ	Yes, if their energies are close
2s and $2p_y$	σ and π	No: zero net overlap
$2p_x$ and $2p_y$	π, but at right angles	No: they are orthogonal
$3d_{xz}$ and $2p_x$	π and π	Yes: a π overlap
$3d_{xy}$ and $3d_{x^2-y^2}$	δ and δ, but rotated 45°	No: orthogonal to each other Both are δ type, yet they cancel; same symmetry label is not enough when the lobes are turned 45°.

Same symmetry about the axis and a matching orientation are both needed for a net overlap.

Unpaired electrons and magnetism10 rows

Species	Electrons	Bond order	Unpaired electrons	Magnetism
$\mathrm{H_2^+}$ , $\mathrm{He_2^+}$	1, 3	0.5	1	Paramagnetic
$\mathrm{Li_2}$	6	1	0	Diamagnetic
$\mathrm{B_2}$	10	1	2	Paramagnetic
$\mathrm{C_2}$	12	2	0	Diamagnetic
$\mathrm{C_2^-}$ , $\mathrm{N_2^+}$	13	2.5	1	Paramagnetic
$\mathrm{N_2}$ , CO, $\mathrm{CN^-}$ , $\mathrm{NO^+}$	14	3	0	Diamagnetic
$\mathrm{N_2^-}$ , $\mathrm{O_2^+}$ , NO	15	2.5	1	Paramagnetic
$\mathrm{O_2}$ , $\mathrm{N_2^{2-}}$	16	2	2	Paramagnetic
$\mathrm{O_2^-}$	17	1.5	1	Paramagnetic
$\mathrm{O_2^{2-}}$ , $\mathrm{F_2}$	18	1	0	Diamagnetic O₂²⁻ has 10 electrons in bonding orbitals and 8 in antibonding ones.

Species with the same electron count have the same bond order and the same number of unpaired electrons.

Watch out for (6)

Bonding π density is not low above the axis
→ Combining atomic orbitals (LCAO)
Maximum overlap, not minimum
→ Combining atomic orbitals (LCAO)
Count every electron, core included, or none
→ Bond order from the MO diagram
A bond order of zero means no molecule
→ Bond order from the MO diagram
O₂⁺ and O₂⁻ have the same number of unpaired electrons
→ Unpaired electrons and magnetism
N₂²⁻ looks like N₂ but behaves like O₂
→ Unpaired electrons and magnetism

Dipole Moment, Hydrogen Bonding and Intermolecular Forces

Formulas (1)

Dipole moment: size and direction · Dipole moment $\mu = q \times d \qquad 1\ \text{D} = 10^{-18}\ \text{esu cm}$

Reference tables (2)

Polar or non-polar: when symmetry cancels9 rows

Shape	Net dipole	Examples
Linear AX₂ or AX₂E₃	Zero	$\mathrm{CO_2}$ , $\mathrm{BeF_2}$ , $\mathrm{BeCl_2}$ , $\mathrm{XeF_2}$
Trigonal planar AX₃	Zero	$\mathrm{BF_3}$ , $\mathrm{BCl_3}$ , $\mathrm{SO_3}$
Tetrahedral AX₄	Zero	$\mathrm{CH_4}$ , $\mathrm{CCl_4}$ , $\mathrm{SiF_4}$
Square planar, TBP, octahedral	Zero	$\mathrm{XeF_4}$ , $\mathrm{PCl_5}$ , $\mathrm{SF_6}$
Bent	Non-zero	$\mathrm{H_2O}$ , $\mathrm{H_2S}$ , $\mathrm{SO_2}$
Pyramidal	Non-zero	$\mathrm{NH_3}$ , $\mathrm{NF_3}$ , $\mathrm{PCl_3}$
See-saw, T-shaped, square pyramidal	Non-zero	$\mathrm{SF_4}$ , $\mathrm{ClF_3}$ , $\mathrm{BrF_5}$
Tetrahedral with mixed atoms	Non-zero	$\mathrm{CHCl_3}$ , $\mathrm{CH_2Cl_2}$
Heteronuclear diatomic	Non-zero	HF, HCl, HBr H₂ has zero dipole; HF, with the biggest electronegativity gap, has the largest of the hydrogen halides.

A lone pair on the centre breaks the symmetry unless the lone pairs themselves are placed symmetrically, as in XeF₂ and XeF₄.

Hydrogen bonding and intermolecular forces7 rows

Case	Kind of attraction	Effect
HF	Intermolecular H-bonds, zig-zag chains	The strongest single H-bond; the H sits nearer one F, so the bonds are not symmetrical
Ice, water, water with solute	Intermolecular H-bonds	Most in ice (each molecule bonded four ways), fewer in liquid water, fewer again with impurities
$o$ -Nitrophenol, salicylaldehyde	Intramolecular H-bond	Lower boiling point; steam volatile
$p$ -Nitrophenol, $p$ -hydroxybenzaldehyde	Intermolecular H-bonds	Higher boiling point; not steam volatile
$\mathrm{CH_4 < HCN < NH_3}$	None, weak C–H···N, N–H···N	Order of intermolecular H-bond strength
Noble gases, $\mathrm{CH_4}$	London forces only	Energy $\propto 1/r^6$ ; grows with molecular size
Ar, $\mathrm{CH_4}$ , $\mathrm{H_2O}$ , $\mathrm{C_6H_6}$	Van der Waals constant a (about 1.4, 2.3, 5.5, 18 L² bar mol⁻²)	Larger a means stronger attraction between molecules

H bonded to F, O or N gives a hydrogen bond; where it forms, inside or between molecules, decides the boiling point.

Watch out for (6)

More electronegative F does not mean a bigger dipole
→ Dipole moment: size and direction
The two arrow conventions point opposite ways
→ Dipole moment: size and direction
Polar bonds can give a non-polar molecule
→ Polar or non-polar: when symmetry cancels
Lone pairs do not always make a molecule polar
→ Polar or non-polar: when symmetry cancels
Ortho means intramolecular
→ Hydrogen bonding and intermolecular forces
HF has no intramolecular hydrogen bond
→ Hydrogen bonding and intermolecular forces

PYQ weightage by concept

21 concepts · 176 PYQs — where the marks actually sit, so you know what to drill first

Lewis Structures, Formal Charge and the Octet Rule19 PYQs · 11%

Concept	PYQs	Share
Counting lone pairs and formal charge	8	5%
The octet rule and its three exceptions	6	3%
Lewis acids and Lewis bases	5	3%

Ionic Bonding, Lattice Enthalpy and Fajans' Rules9 PYQs · 5%

Concept	PYQs	Share
Fajans' rules and covalent character	5	3%
The Born-Haber cycle and lattice enthalpy	4	2%

Bond Length, Bond Angle and Resonance15 PYQs · 9%

Concept	PYQs	Share
Bond angles and lone-pair repulsion	6	3%
Bond length and what sets it	5	3%
Resonance and fractional bond order	4	2%

VSEPR Theory and Lone Pairs16 PYQs · 9%

Concept	PYQs	Share
Counting lone pairs on the central atom	10	6%
Where lone pairs sit: equatorial and trans	6	3%

Shapes of Molecules and Ions27 PYQs · 15%

Concept	PYQs	Share
Shapes from the AXE formula	17	10%
Counting species of one shape in a list	10	6%

Hybridisation and Sigma and Pi Bonds27 PYQs · 15%

Concept	PYQs	Share
Hybridisation from the steric number	11	6%
Counting σ and π bonds	8	5%
Hybridisation, orientation and complexes	8	5%

Molecular Orbital Theory36 PYQs · 20%

Concept	PYQs	Share
Bond order from the MO diagram	15	9%
Unpaired electrons and magnetism	13	7%
Combining atomic orbitals (LCAO)	8	5%

Dipole Moment, Hydrogen Bonding and Intermolecular Forces27 PYQs · 15%

Concept	PYQs	Share
Hydrogen bonding and intermolecular forces	11	6%
Polar or non-polar: when symmetry cancels	9	5%
Dipole moment: size and direction	7	4%

Test yourself on Chemical Bonding and Molecular Structure

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Lewis Structures, Formal Charge and the Octet Rule

Ionic Bonding, Lattice Enthalpy and Fajans' Rules

Bond Length, Bond Angle and Resonance

VSEPR Theory and Lone Pairs

Shapes of Molecules and Ions

Hybridisation and Sigma and Pi Bonds

Molecular Orbital Theory

Dipole Moment, Hydrogen Bonding and Intermolecular Forces

Formula & revision sheet

PYQ weightage by concept

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