JEE Mains Chemistry · Chemical Bonding and Molecular Structure

Lewis Structures, Formal Charge and the Octet Rule

A Lewis structure shares out every valence electron as bond pairs and lone pairs; from it you read the lone-pair count, each atom's formal charge and whether the octet rule holds.

All Chemical Bonding and Molecular Structure notes JEE Mains Chemistry strategy

Why this matters

Nineteen PYQs, twelve of them multiple choice, and six from 2026. Eight count valence electrons, lone pairs or formal charges from a Lewis structure; six sort molecules by how they break the octet rule; five test Lewis acids and bases. Three ideas cover the page.

Concept 1 of 3: Counting lone pairs and formal charge

Every valence electron in a Lewis structure is either in a bond or in a lone pair. So once you know the total and the number of bonds, the lone pairs follow by subtraction. Formal charge is bookkeeping: it compares the electrons an atom owns in the structure with the electrons it brought.

Definition

Total valence electrons: add the group valence electrons of every atom, add one for each negative charge, subtract one for each positive charge.
Each bond (single, double or triple counts as 1, 2 or 3 bonds) uses 2 electrons.
Lone pairs in the whole species $= \dfrac{\text{total valence electrons} - 2 \times \text{bonds}}{2}$ .
Formal charge on an atom $= V - L - \tfrac{1}{2}S$ : $V$ valence electrons of the free atom, $L$ its lone-pair electrons, $S$ the electrons in its bonds.
The formal charges add up to the charge on the species.
Read the question: 'lone pairs in the molecule' counts every atom; 'lone pairs on the central atom' counts one atom only.

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

Worked example

Draw the Lewis structure of the nitrate ion,

\mathrm{NO_3^-}

. How many lone pairs does the ion carry, and what is the formal charge on each atom?

Practice this conceptself-check · 4 quick reps

The same idea in a real exam question:

JEE Mains · 2026 · 23 Jan 2026 Shift 1 · Q40Moderate

Example 1 · Chemical Bonding and Molecular Structure · Lewis Structures, Formal Charge and the Octet Rule

Identify the molecule

(X)

with maximum number of lone pairs of electrons (obtained using Lewis dot structure) among

HNO_{3},H_{2}SO_{4},NF_{3}

and

O_{3}

. Choose the correct bond angle made by the central atom of the molecule (X).

Count every atom, not just the centre

Most of a molecule's lone pairs sit on the outer atoms. In

\mathrm{NF_3}

nitrogen has one lone pair but each fluorine has three, so the molecule has 10. Check whether the question asks for the whole molecule or the central atom.

The charge changes the electron count

An anion has extra electrons and a cation has fewer. Forgetting the charge of

\mathrm{NO_2^-}

gives 17 electrons instead of 18, and a half lone pair, which is the sign that something is wrong.

Drill 7 more on counting lone pairs and formal charge

Concept 2 of 3: The octet rule and its three exceptions

Atoms tend to bond until each has eight electrons around it, like a noble gas. Three kinds of molecule break the rule: some centres stop short of eight, some species have an odd number of electrons, and atoms from period 3 onwards can hold more than eight.

Definition

Count the electrons around the central atom: 2 for every bond pair (a double bond counts twice) and 2 for every lone pair.
Incomplete octet: fewer than eight, usually Be (4), B and Al (6).
Odd-electron species: an odd total of valence electrons, so one electron is unpaired and some atom cannot have eight.
Expanded octet: more than eight, possible only for period 3 and heavier atoms, which have d orbitals to use.
Electron-deficient molecules such as $\mathrm{B_2H_6}$ have too few electrons for every bond to be a normal pair; its bridges are three-centre two-electron bonds.

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.

Practice this conceptself-check · 4 quick reps

The same idea in a real exam question:

JEE Mains · 2024 · 8 April 2024 · Q56Moderate

Example 2 · Chemical Bonding and Molecular Structure · Lewis Structures, Formal Charge and the Octet Rule

Number of molecules from the following which are exceptions to octet rule is ______ .

CO_{2},NO_{2},H_{2}SO_{4},BF_{3},CH_{4},SiF_{4},ClO_{2},PCl_{5}

BeF_{2},C_{2}H_{6},CHCl_{3},CBr_{4}

Sulphur acids and oxides are expanded

Drawn with S=O double bonds, S in

\mathrm{H_2SO_4}

and

\mathrm{SO_3}

has 12 electrons and S in

\mathrm{SO_2}

has 10. JEE keys count all three as expanded octets, so do not call them octet-obeying.

Odd electrons mean an exception, even with a small atom

\mathrm{NO}

and

\mathrm{NO_2}

contain only period 2 atoms, yet they break the rule because 11 and 17 cannot be shared out as complete pairs. Check the total before looking at the central atom.

Drill 5 more on the octet rule and its three exceptions

Concept 3 of 3: Lewis acids and Lewis bases

A Lewis acid accepts an electron pair and a Lewis base donates one. So an electron-poor centre with an empty orbital is an acid, and a centre that still holds a lone pair is a base. The Lewis structure tells you which one you have.

Definition

Lewis acid: an incomplete octet or a vacant orbital, such as B in $\mathrm{BF_3}$ ( $sp^2$ , empty p orbital) or Al in $\mathrm{AlCl_3}$ .
Lewis base: a lone pair on the central atom, such as N in $\mathrm{NF_3}$ or $\mathrm{NH_3}$ ( $sp^3$ ), S in $\mathrm{SF_4}$ , Cl in $\mathrm{ClF_3}$ .
A central atom with no lone pair, such as P in $\mathrm{PCl_5}$ , cannot act as a Lewis base.
Boron halides: acid strength $\mathrm{BI_3 > BBr_3 > BCl_3 > BF_3}$ . The small F atom feeds its lone pair back into boron's empty p orbital ( $p\pi$ – $p\pi$ back-bonding), which makes $\mathrm{BF_3}$ the weakest acid.
Water accepting or giving a proton is Brønsted behaviour, not Lewis behaviour.

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.

Practice this conceptself-check · 4 quick reps

The same idea in a real exam question:

JEE Mains · 2026 · 22 Jan 2026 Shift 1 · Q43Moderate

Example 3 · Chemical Bonding and Molecular Structure · Lewis Structures, Formal Charge and the Octet Rule

Two p-block elements X and Y form fluorides of the type

EF_{3}

. The fluoride compound

XF_{3}

is a Lewis acid and

YF_{3}

is a Lewis base. The hybridization of the central atoms of

XF_{3}

and

YF_{3}

respectively are

Electronegativity does not rank the boron halides

F is the most electronegative halogen, so

\mathrm{BF_3}

looks like it should be the strongest acid. It is the weakest, because back-bonding from F fills boron's empty orbital. The order is

\mathrm{BI_3 > BBr_3 > BCl_3 > BF_3}

Amphoteric water is a Brønsted idea

Water acts as an acid with

\mathrm{NH_3}

and as a base with

\mathrm{H_2S}

by passing protons. An assertion that this is explained by the Lewis concept is false.

Drill 4 more on lewis acids and lewis bases

Summary — formulas & gotchas at a glance

A revision cheat-sheet for the formulas and gotchas above. Click any concept name to jump back to its full explanation.

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

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