Chemical Bonding — Class 10 Physical Science
"Atoms bond to become STABLE. The noble gases already ARE stable. Every other atom is trying to become one. This is the OCTET RULE."
1. The Octet Rule
Atoms combine to achieve 8 ELECTRONS in their OUTERMOST SHELL (octet) — the stable configuration of NOBLE GASES. Exceptions: Hydrogen and Helium — stable with 2 electrons (duet). Atoms achieve octet by: TRANSFER of electrons (ionic bond). SHARING of electrons (covalent bond).
2. Ionic (Electrovalent) Bond — Electron TRANSFER
Metal (low ionisation energy) → LOSES electron(s) → CATION (+). Non-metal (high electron affinity) → GAINS electron(s) → ANION (−). Electrostatic ATTRACTION between opposite ions forms the bond.
Example — NaCl: Na (2,8,1) → loses 1e⁻ → Na⁺ (2,8) — neon configuration. Cl (2,8,7) → gains 1e⁻ → Cl⁻ (2,8,8) — argon configuration. Na⁺ + Cl⁻ → NaCl.
Properties of Ionic Compounds: HIGH melting and boiling points (strong electrostatic forces throughout crystal lattice). CONDUCT electricity when MOLTEN or in AQUEOUS SOLUTION (ions are FREE to move). Do NOT conduct in SOLID state (ions locked in lattice). BRITTLE (crystal layers shift → like charges repel → shatter). Usually SOLUBLE in water. 'NaCl does NOT conduct electricity as a SOLID. It must be MOLTEN (melted) or DISSOLVED for the ions to move freely.'
3. Covalent Bond — Electron SHARING
Two non-metals SHARE valence electrons. Each atom counts the shared electrons toward its octet. Single bond: 1 shared pair (H−H, Cl−Cl). Double bond: 2 shared pairs (O=O). Triple bond: 3 shared pairs (N≡N).
Properties of Covalent Compounds: LOW melting and boiling points (weak INTERMOLECULAR forces — molecules are separate, not in a lattice). Do NOT conduct electricity (no ions, no free electrons). Usually INSOLUBLE in water. Soluble in ORGANIC solvents (alcohol, acetone).
4. Lewis Dot Structures
Show VALENCE ELECTRONS as DOTS around the element symbol. Electrons are paired. For compounds: show shared pairs as lines or dot pairs between atoms.
5. Common Mistakes
- 'NaCl molecules exist independently' — NaCl is an IONIC COMPOUND — a giant CRYSTAL LATTICE. Individual 'NaCl molecules' don't exist in the solid. The formula NaCl represents the simplest RATIO (1:1), not a discrete molecule.
- 'Covalent bonds are weaker than ionic' — The INTRA-molecular covalent bond (within a molecule) is very STRONG. The INTER-molecular forces (between molecules) are WEAK — which is why covalent compounds have LOW melting points.
6. AP SSC Exam Focus
| Topic | Marks |
|---|---|
| Ionic bond formation (NaCl, MgO) | 3-4 |
| Ionic vs Covalent properties | 3-4 |
| Lewis dot structures | 2-3 |
7. Covalent Bonding — Deep Dive
Types of Covalent Bonds
| Bond Type | Shared Pairs | Example | Lewis Structure |
|---|---|---|---|
| Single Bond | 1 pair | H−H, H−Cl, Cl−Cl | H:H or H−H |
| Double Bond | 2 pairs | O=O, C=O | :O::O: or O=O |
| Triple Bond | 3 pairs | N≡N, C≡H | :N:::N: or N≡N |
Example — Formation of H₂O: Oxygen atom: 2, 6 (needs 2 more electrons for octet) Each Hydrogen atom: 1 (needs 1 more electron for duet) Oxygen shares 1 electron with EACH hydrogen atom, forming TWO single covalent bonds. Lewis structure: H−Ö−H (with 2 LONE PAIRS on oxygen — 4 unshared electrons)
'In H₂O, oxygen has TWO bond pairs and TWO lone pairs. The lone pairs cause the bent shape. This is why water is a BENT molecule, not linear.'
Example — Formation of CO₂: Carbon atom: 2, 4 (needs 4 more electrons for octet) Each Oxygen atom: 2, 6 (needs 2 more for octet) Carbon forms a DOUBLE bond with EACH oxygen. Structure: O=C=O 'Each double bond has 4 shared electrons — 2 from carbon and 2 from oxygen.'
Covalent Bonding in Carbon Compounds
Carbon has 4 valence electrons. It forms 4 COVALENT bonds to complete octet. This is the basis of ORGANIC CHEMISTRY (covered in detail in Class 10 Carbon and its Compounds chapter). 'Methane CH₄ has 4 single bonds. Ethene C₂H₄ has one double bond between carbons. Ethyne C₂H₂ has one triple bond between carbons.'
8. Polarity in Covalent Bonds
A covalent bond can be:
| Type | Description | Example |
|---|---|---|
| Non-polar Covalent | Electrons shared EQUALLY. Same atoms or similar electronegativity. | H−H, Cl−Cl, C−H |
| Polar Covalent | Electrons pulled MORE toward one atom. Unequal sharing. | H−Cl, H−O, C−Cl |
Electronegativity Difference determines polarity:
| ΔEN | Bond Type |
|---|---|
| 0 to 0.4 | Non-polar covalent |
| 0.5 to 1.7 | Polar covalent |
| > 1.7 | Ionic (or very polar covalent) |
'Polar covalent bonds create partial charges: δ+ on the less electronegative atom, δ− on the more electronegative atom. In HCl, H has δ+ and Cl has δ−.'
9. Metallic Bonding
Metallic Bond: The electrostatic attraction between a LATTICE of positive metal ions and a SEA of DELOCALISED electrons.
Properties Explained by Metallic Bonding
| Property | Explanation |
|---|---|
| Good conductors of electricity | Delocalised electrons are FREE to move and carry charge |
| Good conductors of heat | Delocalised electrons transfer kinetic energy RAPIDLY |
| Malleable and ductile | Layers of ions can SLIDE over each other without breaking bonds — the electron sea holds them together |
| High melting points | Strong electrostatic attraction between ions and electron sea |
| Lustre (shiny) | Free electrons reflect light |
'Metals are NOT giant molecules with covalent bonds. They are a lattice of positive ions SURROUNDED by a sea of mobile electrons. This is why you can hammer gold into THIN FOIL without shattering it.'
10. Comparison of Bond Types — Summary Table
| Property | Ionic | Covalent | Metallic |
|---|---|---|---|
| Bond formed by | Electron TRANSFER | Electron SHARING | Electron SEA |
| Between | Metal + Non-metal | Non-metal + Non-metal | Metal atoms |
| Melting/Boiling Point | HIGH | LOW | High (varies) |
| Electrical conductivity | In molten/aqueous solution | Does NOT conduct | CONDUCTS (solid/liquid) |
| Solubility in water | Usually SOLUBLE | Usually INSOLUBLE | Insoluble |
| Solubility in organic solvents | Insoluble | SOLUBLE | Insoluble |
| Structure | Crystal lattice | Discrete molecules | Lattice of ions in electron sea |
| Example | NaCl, MgO, CaCl₂ | H₂O, CO₂, CH₄, C₆H₁₂O₆ | Fe, Cu, Au, Al |
11. Lewis Dot Structures — More Examples
Step-by-Step Method for Drawing Lewis Structures
- Count TOTAL valence electrons.
- Identify the CENTRAL atom (usually the least electronegative).
- Connect atoms with SINGLE bonds (2 electrons each).
- DISTRIBUTE remaining electrons as LONE PAIRS to satisfy octet (8 for main group, 2 for H).
- If octet is not complete, form DOUBLE or TRIPLE bonds.
Example — NH₃ (Ammonia): N: 5 valence electrons, H: 1 each × 3 = 3. Total = 8 valence electrons. Structure: H−N−H with one lone pair on N. H
Example — CH₄ (Methane): C: 4, H: 1×4 = 4. Total = 8. Structure: Tetrahedral with C in centre, 4 H atoms at corners. All single bonds, NO lone pairs.
Example — C₂H₄ (Ethene): C: 4×2 = 8, H: 1×4 = 4. Total = 12. H₂C=CH₂ — Double bond between carbon atoms.
12. Common Mistakes — Expanded
| Mistake | Correction |
|---|---|
| 'Ionic compounds have molecules' | Ionic compounds have CRYSTAL LATTICES. No discrete molecules. |
| 'Covalent bonds are weak' | INTRAmolecular covalent bonds are STRONG. INTERmolecular forces are weak. |
| 'All metals are ionic' | Metals have METALLIC bonding — a sea of delocalised electrons. |
| 'NaCl conducts electricity as solid' | Only when MOLTEN or DISSOLVED. Ions must be FREE to move. |
| 'Water has a linear shape' | Water is BENT due to lone pair repulsion (104.5°). |
| 'Double bond is twice as strong as single' | A double bond is STRONGER than a single bond but NOT twice — about 1.5×. |
13. AP SSC Exam Focus — Expanded
| Topic | Marks | Frequency |
|---|---|---|
| Ionic bond formation (NaCl, MgO, CaCl₂) | 3-4 | Very High |
| Ionic vs Covalent properties | 3-4 | Very High |
| Lewis dot structures | 2-3 | High |
| Metallic bonding | 2-3 | Moderate |
| Polarity of bonds | 2 | Moderate |
| Types of covalent bonds (single/double/triple) | 2-3 | High |
14. Self-Test Questions
- Explain the formation of MgO showing electron transfer. Show the Lewis structures.
- Why do ionic compounds have HIGH melting points?
- Draw the Lewis dot structure of CH₄, NH₃, and H₂O.
- Why does NaCl conduct electricity when molten but NOT as a solid?
- Differentiate between ionic, covalent, and metallic bonding.
- What is a polar covalent bond? Give an example.
- Explain why covalent compounds are generally INSOLUBLE in water but SOLUBLE in organic solvents.
- Draw the electron dot structure of N₂ (nitrogen gas showing triple bond).
- Why are metals malleable but ionic compounds are brittle?
- Arrange NaCl, CCl₄, and Fe in increasing order of melting point.
Answers: 1) Mg loses 2e⁻ → Mg²⁺, O gains 2e⁻ → O²⁻, Mg²⁺O²⁻. 2) Strong electrostatic forces throughout the crystal lattice require LARGE energy to break. 3) CH₄: tetrahedral, NH₃: trigonal pyramidal with 1 lone pair, H₂O: bent with 2 lone pairs. 4) Solid state: ions locked in lattice. Molten: ions FREE to move. 5) See comparison table above. 6) Unequal sharing due to electronegativity difference — e.g., H−Cl. 7) Covalent compounds are non-polar. Water is POLAR. 'Like dissolves like' — non-polar dissolves in non-polar (organic solvents). 8) N≡N: with 3 shared pairs and 1 lone pair on each N. 9) Metals: layers SLIDE (electron sea holds). Ionic: layers SHIFT → like charges REPEL → shatter. 10) CCl₄ < Fe < NaCl (covalent < metallic < ionic).
