⚛️ Chemical Bonding – Full Breakdown

ICSE Class 10 Chemistry June 2026

        Chemical Bonding is one of those chapters that feels abstract at first — you can't see bonds, you just have to trust they exist. But once you understand the "why" behind each bond type, it becomes the most scoring chapter. I'll explain everything from scratch: why atoms bother bonding at all, how each bond forms, the exact conditions required, and the kind of reasoning questions ICSE loves. Written like a 93.4% scorer explaining to a friend the night before boards.
    

1) Why Do Atoms Bond? — The Core Reason

Look at the noble gases — He (2), Ne (2,8), Ar (2,8,8). They're perfectly happy. They don't react, don't bond, don't care. Why? Because they have a complete outermost shell (duplet for He, octet for others).

Every other atom wants to achieve this same stable electronic configuration. That's the only reason bonding happens. Atoms are selfish — they want to become stable like noble gases. The driving force is always: lower energy = more stability.

Three ways to get there:

Give away electrons → Ionic bond (metal + non-metal)
Share electrons → Covalent bond (non-metal + non-metal)
Share but one side provides both electrons → Coordinate bond (special covalent)

📖 Out of syllabus but important: The octet rule works for most elements of the 2nd and 3rd periods, but there are exceptions. H and Li are happy with 2 (duplet). Elements in period 3 and beyond can have expanded octets (like S in H₂SO₄ has 12 electrons around it, P in PCl₅ has 10). This is because they have d-orbitals available. Don't mention this in ICSE exam though — stick to octet rule.

2) Ionic (Electrovalent) Bonding

What is it?

Complete transfer of electrons from one atom to another. One becomes positive (cation), the other becomes negative (anion). They stick together because of electrostatic force of attraction between opposite charges.

Conditions Required (very important for reasoning Qs)

Low ionization energy of the metal — it should be able to lose electrons easily (Group 1 & 2 metals are perfect)
High electron affinity of the non-metal — it should readily accept electrons (Halogens and Oxygen are perfect)
Large difference in electronegativity between the two atoms (typically > 1.7 on Pauling scale)
The metal should have 1, 2 or 3 valence electrons (to give away)
The non-metal should have 5, 6 or 7 valence electrons (needs 1-3 to complete octet)

Example — NaCl (Sodium Chloride)

Na (2,8,1) — has 1 valence electron. Cl (2,8,7) — needs 1 electron for octet.

Na loses its 1 electron → becomes Na⁺ (2,8) — stable Ne configuration ✅
Cl gains that electron → becomes Cl⁻ (2,8,8) — stable Ar configuration ✅

Na⁺ and Cl⁻ attract each other → Ionic bond.

Na (2,8,1) ╲ ╲ loses 1 e⁻ ╲ ╰─→ Na⁺ (2,8) [Ne] ╱ Cl (2,8,7) ╱ gains 1 e⁻ ╱ Electrostatic attraction ↔ NaCl lattice

Why NaCl forms a lattice and not just one pair?

One Na⁺ doesn't just stick to one Cl⁻. Each Na⁺ is surrounded by multiple Cl⁻ ions and vice versa. They arrange in a 3D crystal lattice because that's the most energy-efficient arrangement. This is why ionic compounds are solids at room temperature.

Other Examples You Must Know

MgO: Mg (2,8,2) loses 2 → Mg²⁺ (2,8). O (2,6) gains 2 → O²⁻ (2,8). Mg²⁺ + O²⁻ → MgO
CaCl₂: Ca (2,8,8,2) loses 2 → Ca²⁺. Two Cl atoms each gain 1 → 2Cl⁻. Ca²⁺ + 2Cl⁻ → CaCl₂
K₂O: Two K atoms each lose 1 → 2K⁺. O gains 2 → O²⁻. 2K⁺ + O²⁻ → K₂O
MgCl₂: Mg²⁺ + 2Cl⁻ → MgCl₂
Al₂O₃: 2Al³⁺ + 3O²⁻ → Al₂O₃

⚠️ Common board mistake: Students write NaCl as "Na⁺ + Cl⁻ → NaCl" (correct) but then say "NaCl is a molecule" — WRONG. Ionic compounds don't form molecules. They form crystal lattices. NaCl is an ionic compound, not a molecule. The formula unit NaCl tells you the ratio, not a discrete molecule. This exact point comes in 1-mark MCQs and 2-mark reasoning.

Electrovalency

The number of electrons lost or gained by an atom to form an ionic bond. Also called ionic valency.

Na → Na⁺ + 1e⁻ → electrovalency = 1 (positive)
Mg → Mg²⁺ + 2e⁻ → electrovalency = 2 (positive)
Cl + 1e⁻ → Cl⁻ → electrovalency = 1 (negative)
O + 2e⁻ → O²⁻ → electrovalency = 2 (negative)

Properties of Ionic Compounds

Property	Explanation
High melting & boiling points	Strong electrostatic forces between ions → lots of energy needed to break the lattice
Soluble in water	Water is polar. It can pull the lattice apart by surrounding each ion (hydration). Not always — depends on lattice energy vs hydration energy
Insoluble in organic solvents	Non-polar solvents can't separate charged ions
Conduct electricity in molten/aqueous state	Ions are free to move. In solid state, ions are fixed in lattice — no movement → no conduction
Hard but brittle	Lattice is rigid. If you apply force, like layers shift → same charges repel → crystal shatters
Non-directional	Electrostatic attraction is in ALL directions (spherical charge), not a specific direction like covalent bonds

3) Covalent Bonding

What is it?

Two atoms share electrons instead of transferring. Each contributes one electron to form a shared pair. Both atoms "count" this shared pair towards their octet.

Conditions Required

Between two non-metals (or non-metal + same non-metal)
Small difference in electronegativity (typically < 1.7 on Pauling scale)
Both atoms have high ionization energy — they don't want to lose electrons, so they share instead
Both atoms need electrons to complete their octet

Types of Covalent Bonds

a) Single Covalent Bond (shared pair = 1)

H₂: H (1) + H (1) → H:H (each gets duplet of He) ✅
Cl₂: Cl (2,8,7) + Cl (2,8,7) → Cl-Cl (each gets 8e⁻) ✅
H₂O: O (2,6) shares 1 each with two H atoms → H-O-H (O gets 8, H get 2) ✅
CH₄: C (2,4) shares 1 with each of 4 H → C gets 8, each H gets 2 ✅
NH₃: N (2,5) shares 1 with each of 3 H → N gets 8, each H gets 2 ✅

b) Double Covalent Bond (shared pair = 2)

O₂: O=O. Each O has 6 valence electrons. They share 2 pairs (4 electrons). Each O gets 8e⁻. ✅
CO₂: O=C=O. C (4e⁻) shares 2 pairs with each O. Each O shares 2 pairs with C. All get octet.

c) Triple Covalent Bond (shared pair = 3)

N₂: N≡N. Each N has 5e⁻. They share 3 pairs (6 electrons). Each N gets 8e⁻. ✅

Single bond (H₂): H — H (1 shared pair) Double bond (O₂): O ═ O (2 shared pairs) Triple bond (N₂): N ≡ N (3 shared pairs) Lewis dot structure of N₂: :N ≡ N: Each N : 5 dots initially. 3 shared + 2 lone = 8. ✅

⚠️ Board exam note: ICSE ALWAYS asks you to draw the electron dot structure (Lewis structure) of at least one covalent compound. The most common ones: H₂O, NH₃, CH₄, CO₂, N₂, Cl₂, O₂, C₂H₄ (ethene), C₂H₂ (ethyne). Know all of them. I'd say draw them 2-3 times each before the exam so it's muscle memory.

Comparison: Single vs Double vs Triple Bond

Property	Single (C–C)	Double (C=C)	Triple (C≡C)
Bond length	Longest (~154 pm)	Medium (~134 pm)	Shortest (~120 pm)
Bond strength	Weakest (~348 kJ/mol)	Medium (~614 kJ/mol)	Strongest (~839 kJ/mol)
Reactivity	Least reactive	More reactive (π bond breaks)	Most reactive (two π bonds)

Polar vs Non-Polar Covalent Bonds

This is where electronegativity difference matters WITHIN covalent bonds:

Non-Polar Covalent Bond

Electrons shared equally
Same atoms bonded (H₂, O₂, N₂, Cl₂) or atoms with same electronegativity
No charge separation

Polar Covalent Bond

Electrons shared unequally
One atom is more electronegative → it pulls the shared pair closer → gets a slight negative charge (δ⁻), the other gets slight positive (δ⁺)
Examples: H-Cl (Cl more EN), H₂O (O more EN), NH₃ (N more EN)
The greater the electronegativity difference (within 0.4 to 1.7 range), the more polar the bond

Non-polar (Cl₂): Cl — Cl (δ⁻ 0 δ⁺, no separation) Polar (HCl): H — Cl δ⁺ δ⁻ Water (H₂O): δ⁺ H — O — H ↓ δ⁺ δ⁻ (Oxygen pulls electrons)

📖 Out of syllabus: Polarity of a molecule depends on BOTH bond polarity AND molecular shape. CO₂ has polar bonds (C=O) but is non-polar because it's linear — the dipoles cancel out. H₂O has polar bonds AND is bent → net dipole → polar molecule. This is why H₂O dissolves ionic compounds and CO₂ doesn't. ICSE Class 10 doesn't go into molecular geometry in this depth, but knowing this gives you an edge in reasoning questions.

Covalency

The number of electrons an atom shares with other atoms. Basically, the number of covalent bonds it forms.

H → covalency = 1 (forms 1 bond)
O → covalency = 2 (forms 2 bonds, like in H₂O)
N → covalency = 3 (forms 3 bonds, like in NH₃)
C → covalency = 4 (forms 4 bonds, like in CH₄)
Cl → covalency = 1 (forms 1 bond, like in Cl₂ or HCl)

Variable covalency: Some elements show more than one covalency:

P can have covalency 3 (PCl₃) or 5 (PCl₅) — because it has vacant d-orbitals
S can have covalency 2 (H₂S), 4 (SO₂), or 6 (SO₃, H₂SO₄)
N normally has covalency 3, but can have 4 in coordinate bond (NH₄⁺)

⚠️ Very common confusion: Covalency ≠ number of shared electrons. It's the number of covalent bonds. One shared pair = 1 bond = covalency 1. So in N≡N, each N has covalency 3 (three bonds), not 6 (six shared electrons).

4) Coordinate (Dative) Bonding

What is it?

A special type of covalent bond where both shared electrons come from the same atom. The other atom just accepts them. Once formed, a coordinate bond is indistinguishable from a normal covalent bond.

Conditions Required

One atom (donor) must have a lone pair of electrons (unshared pair)
The other atom (acceptor) must be electron-deficient — it needs 2 more electrons to complete its octet
The acceptor should have an empty orbital to accommodate the lone pair

Key Examples for ICSE

a) Ammonium Ion (NH₄⁺)

NH₃ + H⁺ → NH₄⁺

N in NH₃ has a lone pair (it has 5 valence e⁻, uses 3 for covalent bonds with H, 2 remain as lone pair). H⁺ has no electrons — it's just a proton. N donates its lone pair to H⁺. The result: all 4 N-H bonds look identical (3 covalent + 1 coordinate = all same).

H ↘ H — N — H ← coordinate bond (both e⁻ from N) ↙ H (NH₄⁺)

b) Hydronium Ion (H₃O⁺)

H₂O + H⁺ → H₃O⁺

O in H₂O has 2 lone pairs. It donates one to H⁺. Result: 2 normal covalent O-H + 1 coordinate O-H → all equivalent in the ion.

c) Ozone (O₃)

The central O forms one normal covalent bond (sharing one e⁻ each) with one O, and one coordinate bond (donates both e⁻) with the other O. The structure is resonance hybrid — both O-O bonds are identical in reality.

How to Identify Coordinate Bonds in Exams

Look for these patterns:

A central atom with lone pair donating to an electron-deficient species
Formation of NH₄⁺, H₃O⁺, BF₃·NH₃ (B in BF₃ is electron-deficient, N donates lone pair)
In dot structures, coordinate bond is usually shown with an arrow (→) pointing from donor to acceptor

🧠 From a 93.4% scorer: ICSE loves asking the difference between covalent and coordinate bonds in 2-mark questions. Your answer: "In a covalent bond, each atom contributes one electron to the shared pair. In a coordinate bond, both electrons come from the same atom (donor) while the acceptor only provides an empty orbital." That's it. Easy 2 marks.

5) Comparison of All Three Bond Types

Feature	Ionic	Covalent	Coordinate
Electron transfer/sharing	Complete transfer	Equal sharing (each contributes 1)	One contributes both, other contributes none
Between	Metal + Non-metal	Non-metal + Non-metal	Donor (lone pair) + Acceptor (deficient)
EN difference	Large (>1.7)	Small (<1.7)	Varies
Nature	Non-directional (all directions)	Directional (specific angles)	Directional (like covalent)
Melting point	High	Low	Similar to covalent
Electrical conductivity	Conducts when molten/aqueous	Doesn't conduct (except polar in water)	Varies
Solubility in water	Soluble (generally)	Insoluble (most)	Varies

6) Properties of Covalent Compounds

Property	Explanation
Low melting & boiling points	Weak intermolecular forces (van der Waals). Little energy needed to separate molecules.
Generally insoluble in water	Most covalent compounds are non-polar. Water is polar — "like dissolves like". Non-polar solutes don't dissolve in polar solvents.
Soluble in organic solvents	Organic solvents (benzene, CCl₄, alcohol) are also non-polar → dissolve covalent compounds.
Don't conduct electricity	No free ions or electrons. Exception: polar covalent compounds in water (like HCl) can dissociate into ions → conduct.
Directional bonds	Bonds are at specific angles — determines molecular shape (e.g., H₂O bent at 104.5°, NH₃ pyramidal at 107°).

7) Sigma (σ) and Pi (π) Bonds — Beyond ICSE but Useful

📖 This is officially Class 11+ content, but I'm including it because it explains WHY double/triple bonds behave differently. Skip if you don't want overload, but read it if you actually want to understand.

A single bond = 1 σ bond (head-on overlap of orbitals).

A double bond = 1 σ + 1 π bond (σ from head-on, π from sideways overlap of p-orbitals).

A triple bond = 1 σ + 2 π bonds.

π bonds are weaker than σ bonds because sideways overlap is less effective. That's why double and triple bonds are more reactive — the π bonds break first. That's also why C=C bonds undergo addition reactions (π bond breaks open).

8) Top ICSE-Style Reasoning Questions

Q1: Why does NaCl have a high melting point while HCl has a low melting point?

Answer: NaCl is ionic — strong electrostatic forces between Na⁺ and Cl⁻ in a crystal lattice. Requires high energy to overcome. HCl is covalent — weak intermolecular forces (dipole-dipole and van der Waals) between molecules. Easily overcome with little energy.

Q2: Why does NaCl conduct electricity in molten state but not in solid state?

Answer: In solid state, ions are fixed in the lattice and can't move. No charge flow → no conduction. In molten state, the lattice breaks and ions are free to move → they can carry charge.

Q3: What type of bond is formed between two non-metals? Why?

Answer: Covalent bond. Both non-metals have high ionization energy — they don't want to lose electrons. Instead, they share electrons to achieve octet.

Q4: Why is NH₃ a covalent compound but NH₄Cl is ionic?

Answer: NH₃ has covalent bonds between N and H (sharing). NH₄Cl contains NH₄⁺ (ammonium ion) and Cl⁻ (chloride ion). The bond between NH₄⁺ and Cl⁻ is ionic (electrostatic attraction). But within NH₄⁺, the N-H bonds are covalent (3 normal + 1 coordinate).

Q5: Distinguish between ionic and covalent compounds based on melting point and conductivity.

Answer: Table format — Ionic: high MP, conducts when molten/aqueous. Covalent: low MP, generally doesn't conduct.

Q6: Why does oxygen form a double bond but chlorine forms a single bond?

Answer: O has 6 valence electrons — needs 2 more for octet. So it shares 2 pairs (double bond) with another O. Cl has 7 valence electrons — needs only 1 more for octet. So it shares 1 pair (single bond) with another Cl.

⚠️ My final advice from actual board exam experience:
1. Know how to draw electron dot structures of at least 8-10 compounds — they repeat every year.
2. The "differentiate between ionic and covalent" table is a free 4-5 marks — memorise 4-5 points.
3. Coordinate bond questions usually revolve around NH₄⁺ and H₃O⁺ — learn them thoroughly.
4. For "why" questions, always link back to the octet rule / stability. That's the root cause of everything in this chapter.