Introduction

Chemical bonding explains why atoms join together to form molecules and compounds, and why those substances behave the way they do — their melting points, conductivity, shape, and reactivity. It's a core topic across JEE, NEET, KCET and Boards, and it underpins almost every other chemistry topic you'll study.

Why atoms bond

Atoms bond to achieve a more stable, lower-energy electron configuration — often (though not always) resembling the nearest noble gas, a pattern known as the octet rule. Bonding can happen by transferring electrons, sharing electrons, or through a "sea" of delocalised electrons.

Types of chemical bonds

Ionic bond Formed by the complete transfer of one or more electrons from a metal to a non-metal, creating oppositely charged ions held together by electrostatic attraction.
Example: Na loses an electron to become Na⁺; Cl gains it to become Cl⁻ → NaCl.
Covalent bond Formed when two atoms (typically non-metals) share one or more pairs of electrons. Can be single, double or triple depending on how many electron pairs are shared.
Example: H₂ (single bond, one shared pair), O₂ (double bond), N₂ (triple bond).
Metallic bond Metal atoms release some outer electrons into a shared "sea" of delocalised electrons, surrounded by positive metal ions. This explains why metals conduct electricity and are malleable.
Coordinate (dative) bond A special covalent bond where both shared electrons come from the same atom. Example: formation of NH₄⁺ from NH₃ and H⁺.

Molecular shape — VSEPR theory

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shape based on the idea that electron pairs around a central atom arrange themselves to minimise repulsion.

Examples
  • CH₄ — 4 bonding pairs, 0 lone pairs → tetrahedral (109.5°)
  • NH₃ — 3 bonding pairs, 1 lone pair → pyramidal (~107°)
  • H₂O — 2 bonding pairs, 2 lone pairs → bent (~104.5°)
  • CO₂ — 2 bonding regions, 0 lone pairs → linear (180°)

Common mistakes

Watch out for
  • Assuming all bonds between a metal and non-metal are purely ionic — many have significant covalent character (see Fajans' rules).
  • Forgetting lone pairs when predicting shape with VSEPR — lone pairs occupy space and change bond angles even though they aren't "visible" in the final shape.
  • Confusing bond order with bond strength direction — higher bond order (e.g., triple bond) generally means a stronger, shorter bond, not weaker.
  • Mixing up electronegativity difference thresholds — very rough guide: large difference → ionic, small/no difference → covalent, but treat this as a spectrum, not a strict cutoff.

Exam preparation tips

Tip
  • Practice drawing Lewis dot structures — most bonding and shape questions start from getting the structure right.
  • Learn the standard VSEPR shapes table (linear, trigonal planar, tetrahedral, pyramidal, bent, octahedral) with bond angles.
  • For NEET/JEE, be ready to compare bond strength, bond length, and polarity across similar molecules.
  • Boards exams often ask you to explain bonding with real examples (NaCl, H₂O, NH₃, CH₄) — know at least one worked example per bond type.

Summary

Atoms bond to reach a more stable electron configuration, through ionic (electron transfer), covalent (electron sharing), metallic (electron sea), or coordinate bonding. VSEPR theory then predicts the 3D shape of covalent molecules based on electron pair repulsion. Together, bond type and molecular shape explain most of a compound's physical and chemical behaviour.

Confused about a specific molecule's shape or bond type? Ask iBuddie for a step-by-step explanation.

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