Chemical Bonding – Ionic and Covalent Compounds

Chemical Bond: The force of attraction that holds atoms together in a molecule or compound.

Why Bonding Occurs: Atoms bond to achieve stable electronic configuration (octet rule - 8 electrons in valence shell).

Octet Rule: Atoms tend to gain, lose, or share electrons to have 8 electrons in their outermost shell (like noble gases).

Duplet Rule: Hydrogen and helium achieve stability with 2 electrons in valence shell.

Types of Bonds: Ionic bond (electrovalent bond), covalent bond, and coordinate covalent bond.

Valence Electrons: Electrons in outermost shell that participate in bonding.

Definition: Bond formed by complete transfer of electrons from one atom to another, creating ions.

Formation: Metal loses electrons (forms cation), non-metal gains electrons (forms anion). Electrostatic attraction holds them.

Example: NaCl → Na⁺ + Cl⁻. Na loses 1 electron (2,8,1 → 2,8); Cl gains 1 electron (2,8,7 → 2,8,8).

Conditions: Large difference in electronegativity (>1.7), low ionization energy of metal, high electron affinity of non-metal.

Electrovalency: Number of electrons lost or gained. Na has electrovalency +1, Cl has electrovalency -1.

Formula Writing: Total positive charge = Total negative charge. Example: Mg²⁺ and O²⁻ → MgO; Al³⁺ and O²⁻ → Al₂O₃.

Physical State: Solid at room temperature. Hard but brittle.

Melting and Boiling Points: High due to strong electrostatic forces between ions.

Solubility: Soluble in water (polar solvent), insoluble in organic solvents like benzene.

Electrical Conductivity: Conduct electricity in molten state or aqueous solution (due to free mobile ions), not in solid state.

Crystal Structure: Form crystalline lattice with regular arrangement of ions.

Examples: NaCl (common salt), KCl, CaCl₂, MgO, Na₂O, CaO.

Definition: Bond formed by sharing of electrons between atoms.

Formation: Both atoms contribute electrons to form shared electron pairs.

Types: Single bond (1 shared pair), double bond (2 shared pairs), triple bond (3 shared pairs).

Examples: H₂ (H-H, single bond), O₂ (O=O, double bond), N₂ (N≡N, triple bond), CH₄ (4 single bonds).

Covalency: Number of electron pairs shared. H has covalency 1, O has covalency 2, N has covalency 3, C has covalency 4.

Polar and Non-polar: Non-polar (equal sharing - H₂, Cl₂), Polar (unequal sharing - HCl, H₂O).

Physical State: Can be solid, liquid, or gas at room temperature.

Melting and Boiling Points: Generally low due to weak intermolecular forces (except network solids like diamond).

Solubility: Soluble in organic solvents (non-polar), insoluble in water (except some polar covalent compounds).

Electrical Conductivity: Do not conduct electricity in any state (no free ions or electrons).

Molecular Structure: Exist as discrete molecules, not lattice structures.

Examples: H₂O, CO₂, CH₄, NH₃, HCl, sugar, alcohol, wax.

Coordinate Covalent Bond: A covalent bond where both electrons are donated by one atom. Represented by arrow (→).

Example: NH₃ + H⁺ → NH₄⁺ (ammonium ion). N donates both electrons to form bond with H⁺.

Difference Ionic vs Covalent: Ionic (transfer), Covalent (sharing). Ionic (high MP/BP), Covalent (low MP/BP). Ionic (conduct in solution), Covalent (don't conduct).

Electronegativity Difference: >1.7 → Ionic; <1.7 → Covalent; ~0 → Pure covalent.

Formation Pattern: Metal + Non-metal → Ionic; Non-metal + Non-metal → Covalent.

Strength: Triple bond > Double bond > Single bond (more shared electrons = stronger bond).