MHT-CET Chemistry · Teaching notes

Chemical Bonding and Molecular Structure — MHT-CET Chemistry

How atoms join and what shape the result takes — a heavily tested MHT-CET chapter (65 PYQs) that rewards a few master tables (VSEPR shapes, hybridization, molecular-orbital filling) plus one clean formula (bond order). It teaches in five movements, foundations first: (1) ionic and covalent bonding — the octet rule, Lewis structures, Fajans' rules and formal charge; (2) hybridization — the steric-number method (sp, sp², sp³, sp³d, sp³d²) and bond angles; (3) VSEPR theory — predicting molecular geometry from bond pairs and lone pairs; (4) molecular orbital theory — bond order = ½(N_b − N_a), magnetic behaviour and stability; (5) dipole moment, polarity and intermolecular forces — why symmetric molecules are non-polar, and hydrogen bonding. The shape, hybridization, MOT and IMF tables carry the recall load; bond order and dipole moment carry the computation. Every PYQ tagged.

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Subtopic notes

PYQ weightage by concept

21 concepts · 65 PYQs — where the marks actually sit, so you know what to drill first

Ionic and Covalent Bonding, Lewis Structures and Octet Rule13 PYQs · 20%

Concept	PYQs	Share
Fajans' rules — covalent character of an ionic bond	4	6%
Lewis structures, resonance count and electrons around an atom	4	6%
Exceptions to the octet rule	2	3%
Formal charge on an atom in a Lewis structure	2	3%
The octet rule and the three ways atoms bond	1	2%

Hybridization7 PYQs · 11%

Concept	PYQs	Share
The steric-number master table	4	6%
Determining a central atom's hybridization	2	3%
What hybridization is	1	2%
Valence bond theory: sigma and pi bondsfoundation	—	—

VSEPR Theory and Molecular Geometry21 PYQs · 32%

Concept	PYQs	Share
Counting bond pairs and lone pairs on the central atom	8	12%
The master shape table (AXnEm to geometry)	8	12%
The VSEPR premise: electron pairs repel and spread out	3	5%
Bond angles and how lone pairs shrink them	2	3%

Molecular Orbital Theory and Bond Order13 PYQs · 20%

Concept	PYQs	Share
Molecular orbitals and the filling order	4	6%
Bond order from the MO configuration	4	6%
Magnetic behaviour, bond length and stability	3	5%
Bond order and magnetic nature of common species	2	3%

Dipole Moment, Polarity and Intermolecular Forces11 PYQs · 17%

Concept	PYQs	Share
Dipole moment: definition and comparison	3	5%
Types of intermolecular force	3	5%
Hydrogen bonding and boiling point	3	5%
Symmetry: when polar bonds give a zero net dipole	2	3%

Formula & revision sheet

6 formulas · 10 reference tables · 41 gotchas across all subtopics — the exam-eve cheat-sheet

Ionic and Covalent Bonding, Lewis Structures and Octet Rule

Formulas (1)

Formal charge on an atom in a Lewis structure · Formal charge $\text{FC} = V - L - \tfrac{1}{2}\,B$

Reference tables (4)

The octet rule and the three ways atoms bond3 rows

Bond type	How the octet is reached	Formed between	Example
Ionic (electrovalent)	Electrons transferred (lost / gained)	Metal + non-metal	$\text{NaCl}$ , $\text{MgO}$
Covalent	One pair shared, one electron from each atom	Non-metal + non-metal	$\text{H}_2\text{O}$ , $\text{SCl}_2$
Coordinate (dative)	Shared pair donated by one atom only	Donor with a lone pair	$\text{NH}_4^{+}$ , $\text{H}_3\text{O}^{+}$ Once formed, a coordinate bond is identical to any ordinary covalent bond — the label only records where the pair came from.

Transfer = ionic; share = covalent; one-sided share = coordinate.

Fajans' rules — covalent character of an ionic bond4 rows

Factor	Effect on covalent character	Bank example
Smaller cation	More covalent (stronger polariser)	$\text{LiI}$ most covalent among LiCl, LiI, NaCl, NaIQ
Larger anion	More covalent → least ionic	$\text{MI}$ has the lowest ionic character ( $\text{MF}>\text{MCl}>\text{MBr}>\text{MI}$ )Q
Higher cation charge	More covalent	$\text{SnCl}_4$ more covalent than $\text{SnCl}_2$ , $\text{PbCl}_2$ , $\text{SbCl}_3$ Q
Small + highly-charged ions	Highest lattice enthalpy	$\text{BeF}_2$ highest among LiCl, NaCl, $\text{BeF}_2$ , $\text{CaCl}_2$ Q Lattice enthalpy scales with charge density (charge / size), the same driver as polarising power.

Small cation, large anion, high cation charge — all push an ionic bond toward covalent.

Exceptions to the octet rule4 rows

Exception type	Electron count on central atom	Examples
Incomplete octet	Fewer than 8	$\text{BF}_3$ , $\text{BeCl}_2$ , $\text{LiCl}$ Q $\text{LiCl}$ is quoted as incomplete because $\text{Li}^{+}$ has a 2-electron duplet, not an octet.
Expanded octet	More than 8 (uses d-orbitals)	$\text{PCl}_5$ , $\text{SF}_6$ , $\text{H}_2\text{SO}_4$
Odd-electron molecule	Odd total → one unpaired electron	$\text{NO}$ , $\text{NO}_2$ Q
Obeys the octet (for contrast)	Exactly 8	$\text{SCl}_2$ , $\text{H}_2\text{O}$ , $\text{CH}_4$ Q

Fewer than 8 = incomplete; more than 8 = expanded; odd total = odd-electron.

Lewis structures, resonance count and electrons around an atom4 rows

Species / term	Key count or definition	Answer the bank wants
$\text{NO}_2^{-}$ (nitrite)	Double bond can sit on either O	2 resonance (Lewis) structuresQ
Electrons around S in $\text{H}_2\text{SO}_4$	2 single + 2 double bonds = 4 bonds	12 electronsQ
Lewis acid	Electron-pair acceptor	Accepts an electron pair (not 'donates $\text{H}^{+}$ ')Q A Lewis acid need not contain hydrogen — $\text{BF}_3$ is a Lewis acid because boron accepts a lone pair.
Lewis base	Electron-pair donor	Donates an electron pair (e.g. $\text{NH}_3$ )

1 bond = 2 electrons; resonance = the count of equivalent double-bond placements.

Watch out for (10)

A coordinate bond is still a covalent bond
→ The octet rule and the three ways atoms bond
Duplet for H and Li, octet for the rest
→ The octet rule and the three ways atoms bond
Ionic character is the reverse of covalent character
→ Fajans' rules — covalent character of an ionic bond
Charge density, not molar mass, sets lattice enthalpy
→ Fajans' rules — covalent character of an ionic bond
Odd electrons cannot complete an octet
→ Exceptions to the octet rule
Expanded octet needs period-3 (or lower) and d-orbitals
→ Exceptions to the octet rule
Lewis acid = electron-pair acceptor, NOT proton donor
→ Lewis structures, resonance count and electrons around an atom
A double bond is 4 electrons when you total around an atom
→ Lewis structures, resonance count and electrons around an atom
Use half the bonding electrons, not all of them
→ Formal charge on an atom in a Lewis structure
Count lone-pair electrons, not lone pairs
→ Formal charge on an atom in a Lewis structure

Hybridization

Formulas (1)

Determining a central atom's hybridization · Steric number $\text{SN} = (\sigma\text{-bonded atoms}) + (\text{lone pairs on central atom})$

Reference tables (2)

Valence bond theory: sigma and pi bonds3 rows

Bond	Sigma and pi	Example
Single bond	$1\sigma$	$\text{H}-\text{H}$ ; C-C in ethane
Double bond	$1\sigma + 1\pi$	$\text{C}=\text{C}$ in ethene; $\text{O}=\text{O}$
Triple bond	$1\sigma + 2\pi$	$\text{C}\equiv\text{C}$ in ethyne; $\text{N}\equiv\text{N}$

The first bond between two atoms is always a sigma bond; any extra bonds are pi.

The steric-number master table5 rows

Steric number	Hybridization	Geometry	Bond angle	Example
2	$sp$	Linear	$180^\circ$	$BeCl_2$ , $C_2H_2$ Q Acetylene $C_2H_2$ has $sp$ carbons (two atoms + one triple bond that counts as one $\sigma$ ) — the bank's classic $sp$ example.
3	$sp^2$	Trigonal planar	$120^\circ$	$BF_3$ , $C_2H_4$ Q Trigonal planar geometry means $sp^2$ — the answer to 'which hybridisation gives trigonal geometry'.
4	$sp^3$	Tetrahedral	$109.5^\circ$	$CH_4$ , $NH_3$ , $H_2O$
5	$sp^3d$	Trigonal bipyramidal	$90^\circ,\,120^\circ$	$PCl_5$ , $SF_4$ Q $SF_4$ is $sp^3d$ (4 bond pairs + 1 lone pair = SN 5); the lone pair distorts it to a see-saw shape but the hybridization stays $sp^3d$ .
6	$sp^3d^2$	Octahedral	$90^\circ$	$SF_6$ , $XeF_4$ Q $XeF_4$ is $sp^3d^2$ (4 bond pairs + 2 lone pairs = SN 6), square planar — NOT $sp^3$ .

Count the steric number first; the row it lands in gives the hybridization, geometry and angle.

Watch out for (7)

The first bond is always sigma
→ Valence bond theory: sigma and pi bonds
Sigma is stronger than pi; pi locks rotation
→ Valence bond theory: sigma and pi bonds
Hybrids formed = orbitals mixed, not bonds made
→ What hybridization is
Geometry name reports atoms only; SN includes lone pairs
→ The steric-number master table
d-orbitals only appear from steric number 5
→ The steric-number master table
Count lone pairs on the central atom, not just the atoms
→ Determining a central atom's hybridization
A multiple bond is one atom, not two, in the steric count
→ Determining a central atom's hybridization

VSEPR Theory and Molecular Geometry

Formulas (1)

Counting bond pairs and lone pairs on the central atom · Lone pairs on the central atom $\text{lp} = \frac{V - \text{bp}}{2}$

Reference tables (2)

The master shape table (AXnEm to geometry)13 rows

Type (AXnEm)	Bond pairs / Lone pairs	Shape	Ideal bond angle	Example
$\text{AX}_2$	2 / 0	Linear	$180^\circ$	$\text{BeCl}_2$ , $\text{C}_2\text{H}_2$
$\text{AX}_3$	3 / 0	Trigonal planar	$120^\circ$	$\text{BF}_3$
$\text{AX}_2\text{E}$	2 / 1	Bent (angular)	about $119.5^\circ$	$\text{SO}_2$
$\text{AX}_4$	4 / 0	Tetrahedral	$109.5^\circ$	$\text{CH}_4$ , $\text{SiCl}_4$ , $\text{NH}_4^{+}$ Q
$\text{AX}_3\text{E}$	3 / 1	Trigonal pyramidal	$107^\circ$	$\text{NH}_3$
$\text{AX}_2\text{E}_2$	2 / 2	Bent (angular)	$104.5^\circ$	$\text{H}_2\text{O}$ , $\text{SCl}_2$ Q
$\text{AX}_5$	5 / 0	Trigonal bipyramidal	$120^\circ$ and $90^\circ$	$\text{PCl}_5$ Q
$\text{AX}_4\text{E}$	4 / 1	See-saw	$90^\circ$ , $120^\circ$	$\text{SF}_4$ , $\text{TeF}_4$ Q $\text{AB}_4\text{E}$ has a trigonal-bipyramidal parent geometry but a see-saw shape — the bank tests both the type-to-shape and the parent-geometry versions.
$\text{AX}_3\text{E}_2$	3 / 2	T-shaped	about $90^\circ$	$\text{ClF}_3$ , $\text{BrF}_3$ , $\text{ICl}_3$
$\text{AX}_2\text{E}_3$	2 / 3	Linear	$180^\circ$	$\text{XeF}_2$
$\text{AX}_6$	6 / 0	Octahedral	$90^\circ$	$\text{SF}_6$
$\text{AX}_5\text{E}$	5 / 1	Square pyramidal	about $90^\circ$	$\text{BrF}_5$ , $\text{IF}_5$ Q
$\text{AX}_4\text{E}_2$	4 / 2	Square planar	$90^\circ$	$\text{XeF}_4$ Q

Read off the shape from the AXnEm type: count X (bonded atoms) and E (lone pairs), then look up the row.

Bond angles and how lone pairs shrink them5 rows

Molecule	Bond pairs / Lone pairs	Bond angle	Note
$\text{CH}_4$	4 / 0	$109.5^\circ$	Ideal tetrahedral — no lone pair to distort.
$\text{NH}_3$	3 / 1	$107^\circ$	One lone pair shrinks $109.5^\circ$ a little.
$\text{H}_2\text{O}$	2 / 2	$104.5^\circ$	Two lone pairs shrink it further.
$\text{BF}_3$	3 / 0	$120^\circ$	Trigonal planar, no lone pair — full angle.Q
$\text{SO}_2$	2 / 1	about $119.5^\circ$	Bent; one lone pair barely dents the $120^\circ$ parent.Q SO $_2$ is the O–S–O $119.5^\circ$ the bank tests — not $109.5^\circ$ or $180^\circ$ ; its parent is trigonal, not tetrahedral.

Take the ideal angle for the parent geometry, then subtract for each lone pair.

Watch out for (9)

Lone pairs count toward the electron geometry but not the described shape
→ The VSEPR premise: electron pairs repel and spread out
'Regular geometry as expected' means zero lone pairs
→ The VSEPR premise: electron pairs repel and spread out
Count lone pairs on the central atom only
→ Counting bond pairs and lone pairs on the central atom
BF₃ has zero lone pairs — boron is electron-deficient
→ Counting bond pairs and lone pairs on the central atom
H₂O is bent, not linear
→ The master shape table (AXnEm to geometry)
SF₄ is not tetrahedral — it has a lone pair
→ The master shape table (AXnEm to geometry)
Parent geometry versus molecular shape
→ The master shape table (AXnEm to geometry)
SO₂ is 119.5°, not 109.5°
→ Bond angles and how lone pairs shrink them
More lone pairs, smaller angle
→ Bond angles and how lone pairs shrink them

Molecular Orbital Theory and Bond Order

Formulas (2)

Bond order from the MO configuration · Bond order $\text{Bond order} = \tfrac{1}{2}\left(N_b - N_a\right)$
Magnetic behaviour, bond length and stability · Bond order controls length and strength $\text{Bond order} \uparrow \;\Rightarrow\; \text{bond length} \downarrow,\quad \text{bond strength} \uparrow$

Reference tables (1)

Bond order and magnetic nature of common species10 rows

Species	Total electrons	Bond order	Magnetic nature
$\text{H}_2$	2	1	Diamagnetic
$\text{Li}_2$	6	1	Diamagnetic
$\text{N}_2$	14	3	Diamagnetic
$\text{N}_2^+$	13	2.5	Paramagnetic One electron removed from a bonding orbital, so bond order drops to 2.5.
$\text{O}_2$	16	2	Paramagnetic Two unpaired electrons in $\pi^*2p$ — the classic paramagnetic diatomic.
$\text{O}_2^+$	15	2.5	Paramagnetic
$\text{O}_2^-$	17	1.5	Paramagnetic
$\text{F}_2$	18	1	Diamagnetic
$\text{CO}$	14	3	Diamagnetic Isoelectronic with $\text{N}_2$ ; MOT gives bond order 3, not the Lewis double bond.
$\text{NO}$	15	2.5	Paramagnetic Odd-electron molecule: one unpaired electron in a $\pi^*2p$ orbital.

Bond order rises to a maximum of 3 at

\text{N}_2

/CO; paramagnetic species are the ones with an unpaired electron.

Watch out for (7)

Count TOTAL electrons, and adjust for an ion's charge
→ Molecular orbitals and the filling order
Only σ* and π* orbitals count as antibonding
→ Molecular orbitals and the filling order
Ions can have a fractional bond order
→ Bond order from the MO configuration
MOT bond order can differ from the Lewis picture
→ Bond order from the MO configuration
O2 is paramagnetic — the two unpaired electrons
→ Magnetic behaviour, bond length and stability
Higher bond order = shorter bond, not longer
→ Magnetic behaviour, bond length and stability
O2 is paramagnetic even though its bond order is a whole number
→ Bond order and magnetic nature of common species

Dipole Moment, Polarity and Intermolecular Forces

Formulas (1)

Dipole moment: definition and comparison · Dipole moment $\mu = q \times d$

Reference tables (1)

Types of intermolecular force4 rows

Force	Acts between	Strength	Example pair
London / dispersion	Any molecules (even non-polar)	Weakest (grows with size)	CH4 + C2H6 Present in every substance; the ONLY force in non-polar molecules. Largest among HX in HI (biggest, most polarisable).
Dipole–induced dipole (Debye)	One polar + one non-polar molecule	Weak	NH3 + C6H6Q
Dipole–dipole	Two polar molecules	Moderate (bigger dipole → stronger)	HF, HCl (polar HX) Strongest dipole–dipole among the hydrogen halides is HF, because F gives the largest bond dipole.
Hydrogen bonding	H on N/O/F, near a lone pair on N/O/F	Strongest of these	H2O, NH3, HF, alcohols

Polar–polar → dipole–dipole; polar–non-polar → dipole-induced dipole; non-polar only → dispersion; H on N/O/F → hydrogen bond.

Watch out for (8)

Dipole moment is a vector — add directions, not magnitudes
→ Dipole moment: definition and comparison
Bigger electronegativity difference → bigger bond dipole
→ Dipole moment: definition and comparison
Polar bonds do NOT guarantee a polar molecule
→ Symmetry: when polar bonds give a zero net dipole
CHCl3 is polar; CCl4 is not
→ Symmetry: when polar bonds give a zero net dipole
Dipole–dipole vs dispersion point to different HX
→ Types of intermolecular force
Match the force to the pair's polarity
→ Types of intermolecular force
No H on N/O/F means no hydrogen-bond donor
→ Hydrogen bonding and boiling point
H2S does not hydrogen-bond like H2O
→ Hydrogen bonding and boiling point