Explain the formation of a chemical bond.

Exercises · 4.1

Q. Explain the formation of a chemical bond.

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Concept understanding — Lewis Dot Structures

Why do we need Lewis dot structures?

Atoms are held together in molecules by chemical bonds — but what exactly is a bond? In the early 20th century, Gilbert N. Lewis realised that the key lies in the valence electrons (the outermost electrons). He noticed that atoms of noble gases (like Ne, Ar) are extremely stable and unreactive, and they all have 8 electrons in their outermost shell (except helium, which has 2). This led to the octet rule: atoms tend to gain, lose, or share electrons to achieve a full outer shell of 8 electrons (or 2 for hydrogen).

Lewis dot structures are simply a shorthand picture of this idea. They show:

Which atoms are connected to which
How many valence electrons each atom contributes
How those electrons are arranged as bonding pairs (shared) or lone pairs (unshared)

The precise statement

A Lewis dot structure (or electron dot structure) represents the valence electrons of an atom or molecule using dots placed around the element's symbol. Each dot stands for one valence electron. Shared pairs (bonds) are shown as lines, and unshared pairs as pairs of dots.

For a single atom, you write the element symbol and place dots on its four sides (top, bottom, left, right) — up to 8 dots. The order of filling doesn't matter for the final picture, but conventionally you place one dot on each side first, then pair them up.

For example:

Carbon (group 14, 4 valence electrons): $\cdot \text{C} \cdot$ (four single dots)
Oxygen (group 16, 6 valence electrons): $\cdot \ddot{\text{O}} \cdot$ (two single dots and two pairs)

How to draw a Lewis structure for a molecule

Here's the step-by-step method you'll use in exams:

Count total valence electrons — add up valence electrons from all atoms. For ions, add 1 electron for each negative charge, subtract 1 for each positive charge.
Identify the central atom — usually the least electronegative element (not hydrogen or fluorine). Place it in the centre.
Connect atoms with single bonds — each bond uses 2 electrons. Subtract these from your total.
Complete octets of outer atoms — place remaining electrons as lone pairs on terminal atoms (except hydrogen, which only needs 2).
Place leftover electrons on the central atom — if any remain.
If the central atom has fewer than 8 electrons, form multiple bonds — move lone pairs from outer atoms to create double or triple bonds until the central atom has an octet.

Watch out

A common mistake: forgetting that hydrogen only needs 2 electrons (a duet), not 8. Never put more than 2 electrons around H.

A concrete example: water (H₂O)

Total valence electrons: O has 6, each H has 1 → $6 + 1 + 1 = 8$ electrons.
Central atom: oxygen (least electronegative after H).
Connect: O—H bonds (2 bonds × 2 electrons = 4 electrons used).
Remaining: $8 - 4 = 4$ electrons → place as two lone pairs on oxygen.
Check: O has 2 bonds (4 electrons) + 2 lone pairs (4 electrons) = 8. Each H has 1 bond (2 electrons) = 2. Done.

The structure: $\text{H} - \ddot{\text{O}} - \text{H}$

What the structure tells you

Once drawn, a Lewis structure reveals:

Bond order (single, double, triple)
Lone pairs (which affect molecular shape and reactivity)
Formal charge (a bookkeeping tool to check which structure is most stable)
Resonance (when multiple valid structures exist for the same molecule)

Important

Lewis structures are not 3D pictures — they show connectivity and electron arrangement in a flat diagram. Actual molecular shapes are determined by VSEPR theory, which builds on the Lewis structure.

Limitations you should know

The octet rule works beautifully for second-period elements (C, N, O, F) but has exceptions:

Incomplete octet: Be, B, Al (stable with 4 or 6 electrons)
Expanded octet: elements in period 3 and beyond (P, S, Cl) can have more than 8 electrons
Odd-electron molecules: NO, NO₂ (can't give every atom an octet)

For your exams, always check if the central atom is from period 2 or below — that tells you whether expanded octets are possible.

Final takeaway

Lewis dot structures are your first tool for understanding bonding. They translate the abstract idea of electron sharing into a simple visual language. Master the counting and placement steps, and you'll be able to draw any small molecule's structure confidently.

Lewis Dot Structures are taught in the NCERT Class 11 Chemistry chapter on Chemical Bonding and Molecular Structure, matching searches like "Lewis structure: rules and examples" or "chemical bonding important questions class 11 chemistry". Drawing these structures correctly is foundational for JEE Main and NEET chemistry questions on VSEPR shapes, hybridisation, and resonance that build directly on this skill.

A chemical bond forms because atoms achieve lower energy (greater stability) by sharing or transferring electrons, following the octet rule. The key idea: nature prefers lower energy states, and bonding is the mechanism atoms use to get there.

The Core Idea: Why Atoms Bother to Bond

Atoms don't bond because they "like" each other. They bond because a bonded state has lower energy than the isolated atoms. Think of a ball at the top of a hill — it will roll down to the bottom because that's a lower energy position. Similarly, atoms rearrange their electrons to reach a more stable, lower-energy configuration.

The driving force is the octet rule: atoms tend to gain, lose, or share electrons to achieve the electron configuration of the nearest noble gas (8 electrons in their valence shell, except for hydrogen and helium which aim for 2). Noble gases are famously unreactive because their electron shells are completely filled — that's the "bottom of the hill" for most atoms.

The Two Main Ways to Bond

Atoms have two fundamental strategies to achieve a full outer shell:

Transfer electrons (ionic bonding) — one atom gives electrons away, another takes them.
Share electrons (covalent bonding) — atoms pool electrons so both can count them.

Let's walk through each.

1. Ionic Bonding: The Give-and-Take

Consider sodium ( $\text{Na}$ ) and chlorine ( $\text{Cl}$ ).

Sodium has 1 valence electron (configuration $2,8,1$ ). It's much easier to lose that one electron than to gain seven.
Chlorine has 7 valence electrons ( $2,8,7$ ). It needs just one more to complete its octet.

When they meet, sodium transfers its single valence electron to chlorine:

\text{Na} \rightarrow \text{Na}^+ + e^- \quad \text{(sodium becomes a cation)}

\text{Cl} + e^- \rightarrow \text{Cl}^- \quad \text{(chlorine becomes an anion)}

Now both have noble gas configurations: $\text{Na}^+$ has $2,8$ (like neon), and $\text{Cl}^-$ has $2,8,8$ (like argon).

Watch out

A common mistake is thinking the bond is "between" just one Na and one Cl. In reality, the oppositely charged ions arrange themselves in a giant lattice — each $\text{Na}^+$ is surrounded by several $\text{Cl}^-$ ions and vice versa. The formula $\text{NaCl}$ tells you the ratio, not that there are discrete molecules.

The electrostatic attraction between these opposite charges is the ionic bond. The energy released when the lattice forms (lattice energy) is what makes the process favourable overall.

2. Covalent Bonding: The Shared Pool

Now take two chlorine atoms. Both need one electron. Neither will give one away — they'd just become unstable ions. So they share a pair of electrons.

Each chlorine atom contributes one electron to the shared pair:

\text{Cl} \cdot + \cdot \text{Cl} \rightarrow \text{Cl} : \text{Cl}

Now each chlorine "counts" the shared pair as part of its own octet. The shared pair is the covalent bond.

Tip

In Lewis dot structures, each dot represents a valence electron. A shared pair is drawn as a line ( $-$ ). So $\text{Cl} - \text{Cl}$ means a single covalent bond — two shared electrons.

3. The Special Case: Coordinate Bond

Sometimes one atom provides both electrons for the shared pair. This is a coordinate bond (or dative bond). For example, in the ammonium ion $\text{NH}_4^+$ :

Nitrogen has a lone pair (two unshared electrons).
A hydrogen ion $\text{H}^+$ has no electrons — it's just a proton.
Nitrogen "donates" its lone pair to form a bond with $\text{H}^+$ .

The resulting bond is identical to any other N–H bond in the ion, but its origin is different. In Lewis structures, it's often shown as an arrow pointing from donor to acceptor.

The Energy Picture: Why Bonding Releases Energy

When two atoms approach each other:

Attractive forces (nucleus of one atom to electrons of the other) pull them together.
Repulsive forces (nucleus-nucleus and electron-electron) push them apart.

At the bond length (the equilibrium distance), these forces balance. The system's potential energy is at a minimum — lower than when the atoms were far apart. The energy released when the bond forms equals the bond energy.

\text{A-B}

Putting It All Together: The Formation of a Chemical Bond

Two or more atoms approach each other.
Their valence electrons interact — either by transfer (ionic) or sharing (covalent).
The atoms rearrange to achieve stable octets (or duets for H/He).
The system's total energy decreases — the bonded state is more stable than the isolated atoms.
The atoms are held together by electrostatic forces (ionic: ion-ion attraction; covalent: shared electrons attracted to both nuclei).

Important

The fundamental condition for bond formation: the potential energy of the system must decrease. If bonding doesn't lower energy, it won't happen — that's why noble gases don't form bonds under normal conditions.

Summary Table: Ionic vs Covalent vs Coordinate

Feature	Ionic	Covalent	Coordinate
Electron transfer	Complete transfer	Equal sharing	One atom donates both electrons
Bonding particles	Ions (cations & anions)	Atoms	Atoms
Typical between	Metal + Non-metal	Non-metal + Non-metal	Lewis base + Lewis acid
Example	$\text{NaCl}$	$\text{H}_2$ , $\text{O}_2$	$\text{NH}_4^+$

✓Final answer

A chemical bond forms when atoms interact to achieve a lower energy state, typically by transferring or sharing valence electrons to complete their octets, resulting in a stable bonded arrangement.

Same / Similar Concept — real previous-year questions on the same or a closely similar concept, not this exact question.

CBSE 2026Set ANNUAL1 markMCQ
$Q. Lewis dot structure of CO, NO2⁻ and CO3²⁻ are I, II and III respectively (see figure). Which of the above structure(s) is/are wrong?$
(a) Only I
(b) Only II
(c) Only III
(d) None of these
›Reveal solutionSolution

Checking the electron count for CO, $NO_2^-$ and $CO_3^{2-}$ shows each structure described is a standard, chemically valid Lewis dot structure (using the common convention of showing only the overall ionic charge rather than individual formal charges).

I — CO: Total valence electrons $= 4(C) + 6(O) = 10$ . A triple bond uses 6 electrons, leaving 4 electrons = one lone pair on C and one on O, exactly as shown. This is the accepted Lewis structure of carbon monoxide (giving formal charges $-1$ on C, $+1$ on O, consistent with its known triple-bond character).

II — $NO_2^-$ : Total valence electrons $= 5(N) + 2(6)(O) + 1 = 18$ . One N=O double bond + one N–O single bond uses 6 electrons; the remaining 12 electrons distribute as lone pairs to complete octets on both oxygens (and a lone pair on N, which also accounts for the ion's bent shape) — consistent with the standard resonance form of nitrite drawn in the figure.

III — $CO_3^{2-}$ : Total valence electrons $= 4(C) + 3(6)(O) + 2 = 24$ . One C=O double bond + two C–O single bonds use 8 electrons; the remaining 16 electrons complete the octets of all three oxygens, with the overall $2-$ charge shown on the ion — this matches a standard resonance structure of carbonate.

Since each structure satisfies the octet rule and the correct overall electron/charge count for a valid resonance form of that species, none of them is wrong.

✓Final answer
$(d) None of these — I, II and III are each valid Lewis dot structures.$
CBSE 2026Set ANN1 markMCQ
$Q. Identify the incorrect Lewis structure.$
(a) O=C=O (each O with two lone pairs)
(b) N=O (N with a single unpaired electron, O with lone pairs)
(c) N(triple bond)N: (with lone pair on one N)
(d) H-C(triple bond)C-H
›Reveal solutionSolution
N2 must show a lone pair on both nitrogens; the option with a lone pair on only one N is the incorrect Lewis structure. Answer: (c).

Check each structure:

(a) O=C=O with two lone pairs on each oxygen: correct Lewis structure of CO2 (carbon has 4 bonds, each O has 2 lone pairs).

(b) N=O with an unpaired electron on N and lone pairs on O: nitric oxide (NO) is an odd-electron molecule, so this is acceptable.

(c) N(triple bond)N with a lone pair on only one nitrogen: incorrect. In dinitrogen each N has 5 valence electrons; after forming a triple bond each nitrogen still has ONE lone pair, so the correct structure is :N(triple)N:. Showing a lone pair on only one atom leaves the other N electron-deficient - hence incorrect.

(d) H-C(triple bond)C-H: correct structure of ethyne.

✓Final answer
$(c) - N2 must show a lone pair on each nitrogen; the given structure is incorrect.$
CBSE 2025Set ANNUAL1 markMCQ
$Q. Which of the following contains both covalent and ionic bonds?$
(a) CCl4
(b) CaCl2
(c) NH4Cl
(d) H2O
›Reveal solutionSolution

NH4Cl contains both covalent bonds (within the NH4+ ion) and an ionic bond (between NH4+ and Cl-).

In NH4Cl, the four N-H bonds within the ammonium ion (NH4+) are covalent (electron sharing), while the bond between the NH4+ cation and the Cl- anion is ionic (electrostatic attraction). CCl4 and H2O are purely covalent compounds; CaCl2 is purely ionic.

✓Final answer
$(C) NH4Cl.$
CBSE 2025Set ANNUAL1 mark
$Q. Answer in one word/sentence: Draw the Lewis structure for H2S.$
›Reveal solutionSolution
Sulfur (Group 16, 6 valence electrons) forms one single bond to each hydrogen, using 4 of its 6 electrons for 2 lone pairs.

Step 1: Count total valence electrons.

S: 6 valence electrons

2 H: 1 electron each = 2

Total = 8 electrons = 4 electron pairs.

Step 2: Sulfur is the central atom (less electronegative than... well H is bonded terminally); place the 2 H atoms around S and form one single (2-electron) bond to each:

$H-S-H$

This uses 2 bond pairs (4 electrons).

Step 3: The remaining $8 - 4 = 4$ electrons are placed as 2 lone pairs on sulfur (satisfying the octet: 2 bond pairs + 2 lone pairs = 8 electrons around S).

Lewis structure:

$H-\overset{\displaystyle \cdot\cdot}{\underset{\displaystyle \cdot\cdot}{S}}-H$

(two dots above and two dots below the S, representing the two lone pairs)

Because of the two lone pairs on S, the molecule is bent/angular in actual shape (similar to water), though the Lewis structure itself is drawn linearly for simplicity.

✓Final answer
$_2$
CBSE 2023Set ANNUAL1 mark
$Q. Draw the Lewis structure of CO3^2- ion.$
›Reveal solutionSolution
CO3^2- Lewis structure: C double-bonded to one O and single-bonded to two O⁻ atoms, with the negative charge and double-bond character delocalized over all three C–O bonds by resonance.

Total valence electrons in CO3^2-: $C(4) + 3\times O(6) + 2(\text{extra for charge}) = 4+18+2 = 24$ electrons (12 pairs).

One Lewis (contributing) structure: place carbon at the centre bonded to three oxygen atoms.

One C–O bond is a double bond (C=O); that oxygen carries 2 lone pairs.

The other two C–O bonds are single bonds; each of those oxygens carries 3 lone pairs and an overall −1 formal charge.

This uses all 24 electrons: 4 bonding pairs (1 double + 2 single = 4 shared pairs = 8 electrons) + 8 lone pairs on the three oxygens (16 electrons) = 24 electrons. ✓

This structure has a net charge of −2 (each singly-bonded O is −1, the doubly bonded O and C are neutral), matching the ion's charge.

Because any of the three oxygens could equally be the doubly-bonded one, there are three equivalent resonance structures, and the real ion is a hybrid of all three: all three C–O bonds are actually identical, each with a bond order of $4/3$ and a partial charge of $-2/3$ on each oxygen.

✓Final answer
$A central carbon atom double-bonded to one oxygen (2 lone pairs on that O) and single-bonded to two other oxygens, each carrying a -1 charge and 3 lone pairs; the true CO3^2- structure is the resonance hybrid of the three such structures, with the double-bond character and negative charge spread equally over all three C-O bonds.$
CBSE 2023Set ANNUAL1 mark
$Q. Match the column: Column I item 'Covalent Bond' is to be matched with the correct term from Column II:$
(a) Sharing of electron,
(b) increasing,
(c) To find out molecular mass of organic compound,
(d) SF6,
(e) Transition element,
(f) Molarity,
(g) Mulliken,
(h) Heisenberg,
(i) Molality. Which Column II term matches this Column I item?
›Reveal solutionSolution

A covalent bond is defined by the sharing of an electron pair between two atoms (option a), unlike an ionic bond which involves electron transfer.

According to the Lewis theory of bonding, a covalent bond forms when two atoms share one or more pairs of electrons so that both atoms attain a stable (often octet) electron configuration. For example, in H2, each hydrogen atom contributes one electron to the shared pair; in Cl2, each chlorine atom contributes one electron of the shared pair.

This is fundamentally different from an ionic bond, which involves complete transfer of electron(s) from one atom to another, creating oppositely charged ions.

✓Final answer
$(a) Sharing of electron.$

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Q.Explain the formation of a chemical bond.