Chemical Bonding — Telangana (SSC) Class 10 Science Solutions (Free)

Chemical bond: Force of attraction that holds atoms together in molecules or compounds. Ionic bond: Formed by complete transfer of electrons from metal to non-metal. Results in electrostatic attraction between cations and anions. Example: NaCl (Na⁺ Cl⁻) Covalent bond: Formed by sharing of electron pairs between non-metals. Electrons are shared equally (nonpolar) or unequally (polar). Example: Cl₂, H₂O Ionic = transfer, Covalent = sharing

Ionic bond: Electrostatic force between oppositely charged ions formed by electron transfer. Formation of NaCl: Na (11e⁻) → Na⁺ + 1e⁻ (loses 1 valence electron) Cl (17e⁻) + 1e⁻ → Cl⁻ (gains 1 electron) Na⁺ [Cl:⁻] (Cl with octet: 8 electrons, Na with empty M shell = 2 electrons in L shell) Electrostatic attraction between Na⁺ and Cl⁻ forms ionic bond. Result: Solid crystal lattice of alternating Na⁺ and Cl⁻ ions

Covalent bond: Chemical bond formed by sharing of electron pair(s) between non-metal atoms. Single bond: Sharing of 1 electron pair (2 electrons). Example: H—H (H₂) or C—C Double bond: Sharing of 2 electron pairs (4 electrons). Example: O=O (O₂) or C=O in CO₂ Triple bond: Sharing of 3 electron pairs (6 electrons). Example: N≡N (N₂) or C≡C Triple bond is strongest, single bond is weakest

Lewis structure of H₂O: O has 6 valence electrons, each H has 1 valence electron. Total = 6 + 1 + 1 = 8 electrons Structure: H—O—H with 2 lone pairs on O Electron distribution: ..O.. (4 dots for 2 lone pairs) H:O:H O forms 2 single covalent bonds (2 pairs shared with H atoms) O has 2 lone pairs (unshared electrons) Each H achieves duet (2 electrons), O achieves octet (8 electrons)

Polar covalent bond: Covalent bond in which electron pair is shared unequally between atoms. Reason: Atoms have different electronegativities. Electrons shift toward more electronegative atom → partial charges develop. Example: H—Cl Cl is more electronegative than H → electrons shift toward Cl H becomes slightly positive (δ+), Cl becomes slightly negative (δ⁻) Result: H^δ+ — Cl^δ⁻ (polar bond) Other examples: H—O, H—N (all bonds to highly electronegative atoms)

Ordinary covalent bond: Each atom contributes one electron to form shared pair. Example: H—H (each H contributes 1 electron) Coordinate covalent (dative) bond: Both electrons in shared pair come from one atom (donor), other atom accepts (acceptor). Still formed same way but origin is different. Example: NH₄⁺ (ammonium ion) — N donates lone pair to H⁺ Once formed, coordinate bond is identical to ordinary covalent bond. Symbol: → (shows direction from donor to acceptor) but often written as regular line

O₂ formation: Each O has 6 valence electrons, total = 12 electrons in O₂ Lewis structure shows double bond: O=O Bond order = (bonding electrons - antibonding electrons)/2 = (8-4)/2 = 2 Double bond = 2 shared pairs (4 electrons) Each O achieves octet configuration O₂ is paramagnetic (has unpaired electrons in antibonding orbitals) Bond strength is high due to double bond, explains why O₂ is stable molecule