Cupboard

1. Chemical Bonding - Detailed PPT
Batch 1: Introduction & Ionic Bond
(Slides 1–20)

2. Introduction
• Atoms rarely exist independently.
• Molecules form when total energy decreases → stability increases.
• Chemical Bond = Attractive force holding atoms together.
• Examples: H2, O2, Cl2, NaCl.

3. Need for Bonding
• Atoms combine to achieve:
• • Minimum Energy State (more stable).
• • Noble Gas Configuration (octet rule).
• Bond formation = Exothermic process (energy release).

4. Definition of Chemical Bond
• The attractive force that holds atoms together in a stable molecule.
• Involves redistribution of valence electrons.
• Bond formation → energy decreases → stability increases.

5. Lewis Symbols
• Valence electrons are represented as dots around element symbols.
• Examples:
• • H (1 electron): H•
• • O (6 electrons): O••••
• • N (5 electrons): N•••
• • Cl (7 electrons): Cl•••••••

6. Octet Rule
• Atoms tend to achieve 8 electrons in outermost shell (stable configuration).
• Achieved by:
• • Losing electrons (Na → Na+).
• • Gaining electrons (Cl → Cl−).
• • Sharing electrons (H2, O2, N2).

7. Exceptions to Octet Rule
• 1. Electron deficient molecules (BF3, BeCl2).
• 2. Odd-electron molecules (NO2, NO).
• 3. Expanded octet (SF6, PCl5, XeF4).
• 4. Does not explain shapes or stability of molecules.

8. Causes of Chemical Combination
• 1. Tendency to achieve minimum energy.
• • Nucleus of one atom attracts electrons of another.
• • Net attraction → potential energy decreases → bond forms.
• 2. Noble Gas Configuration.
• • Atoms combine to achieve stable electronic configuration.

9. Ionic Bond - Definition
• Formed by complete transfer of electrons.
• Results in positively charged cation & negatively charged anion.
• Held by strong electrostatic attraction.
• Examples: NaCl, MgCl2, CaO.

10. Conditions for Ionic Bond
Formation
• 1. Low Ionization Energy (easy cation formation).
• 2. High Electron Affinity (easy anion formation).
• 3. High Lattice Energy (strong ionic bond stability).
• Favourable: Metals (s-block) + Non-metals (p-block).

11. Example: Formation of NaCl
• Na (1s²2s²2p⁶3s¹) → Na+ (1s²2s²2p⁶) + e−
• Cl (1s²2s²2p⁶3s²3p⁵) + e− → Cl− (1s²2s²2p⁶3s²3p⁶)
• Na+ and Cl− held by electrostatic force.
• Crystal lattice forms with high stability.

12. Born-Haber Cycle (Basics)
• Used to calculate lattice energy of ionic compounds.
• Steps:
• • Sublimation of metal atom.
• • Dissociation of non-metal molecule.
• • Ionization of metal atom.
• • Electron gain by non-metal atom.
• • Combination to form ionic crystal.

13. Lattice Energy
• Energy released when one mole of ionic crystal forms from gaseous ions.
• Greater lattice energy → more stable ionic compound.
• Factors:
• • Higher charge = higher lattice energy.
• • Smaller ionic radius = higher lattice energy.

14. Properties of Ionic Compounds
• • Hard, crystalline solids.
• • High melting & boiling points.
• • Non-conductors in solid state (ions fixed).
• • Conduct electricity in molten/aqueous state.
• • Soluble in polar solvents (like water).

15. Examples of Ionic Compounds
• NaCl – common salt.
• MgO – refractory material.
• CaF2 – fluorite, used in metallurgy.
• Al2O3 – alumina, amphoteric oxide.

16. Comparison: Ionic vs Covalent
Bond (Intro)
• Ionic Bond:
• • Transfer of electrons.
• • Non-directional.
• • Strong electrostatic forces.
• Covalent Bond:
• • Sharing of electrons.
• • Directional.
• • Strong but weaker than ionic.