Chemical Bonding Made Simple: Why Atoms Stick Together
A concept-first guide to ionic, covalent and coordinate bonds, VSEPR shapes, hybridisation and molecular orbital theory for JEE Chemistry.
By the PadhoDost Team · 📖 8 min read · Updated 4 August 2026
Part of JEE (Main + Adv) prep🧠 Atoms are like people at a party
Some atoms are 'givers' who happily hand over spare electrons, some are 'sharers' who hold hands over a shared pair, and a few are 'generous hosts' who supply both electrons of a shared pair. Just as people team up to feel comfortable, atoms bond to reach the stable, contented electron arrangement of a noble gas.
Noble gases like neon and argon almost never react because their outermost shell already holds eight electrons (a full octet), which is extremely stable. Every other atom 'wants' that same stability. It can get there by losing, gaining, or sharing electrons. This drive to complete the octet is the single idea behind almost all chemical bonding.
Three ways to bond
| Feature | Ionic bond | Covalent bond | Coordinate bond |
|---|---|---|---|
| How it forms | Transfer of electrons (metal to non-metal) | Sharing of electrons between atoms | One atom donates both shared electrons |
| Typical between | Metal + non-metal | Two non-metals | Lewis base to Lewis acid |
| Example | NaCl | H2, Cl2, CH4 | NH4+, H3O+ |
| General nature | Solids, high melting point | Gases or liquids, lower melting point | Behaves like a normal covalent bond once formed |
VSEPR: predicting molecular shapes
VSEPR (Valence Shell Electron Pair Repulsion) theory says the electron pairs around a central atom - both bonding pairs and lone pairs - repel each other and spread out as far apart as possible. Count the total number of pairs (the steric number) and the geometry follows automatically. Lone pairs are 'fatter' and repel more strongly than bonding pairs, so they slightly squeeze the bond angles.
| Electron pairs (steric no.) | Hybridisation | Shape | Bond angle | Example |
|---|---|---|---|---|
| 2 | sp | Linear | 180 degrees | BeCl2, CO2 |
| 3 | sp2 | Trigonal planar | 120 degrees | BF3 |
| 4 | sp3 | Tetrahedral | 109.5 degrees | CH4 |
| 5 | sp3d | Trigonal bipyramidal | 120 and 90 degrees | PCl5 |
| 6 | sp3d2 | Octahedral | 90 degrees | SF6 |
📝 Worked example: shape of H2O and bond order of O2
Q: Find the shape of a water molecule (H2O) and the bond order of O2.
H2O: Oxygen has 6 valence electrons - 2 are used to bond the two H atoms, leaving 2 lone pairs.
Steric number = 2 bond pairs + 2 lone pairs = 4, so oxygen is sp3 hybridised (tetrahedral electron geometry).
The 2 lone pairs repel more strongly and squeeze the H-O-H angle from 109.5 degrees down to about 104.5 degrees - giving a bent (angular) shape.
O2 by MOT: bonding electrons Nb = 10 and antibonding electrons Na = 6, so bond order = (10 - 6)/2 = 2 (a double bond).
Two electrons sit unpaired in the pi-star (antibonding) orbitals, so O2 is paramagnetic - it is attracted by a magnet.
Quick recap
- ✓Atoms bond to reach a stable noble-gas electron arrangement (the octet rule).
- ✓Ionic bonds transfer electrons, covalent bonds share them, coordinate bonds share a pair donated entirely by one atom.
- ✓VSEPR: count bond pairs + lone pairs (steric number) to predict shape; lone pairs shrink bond angles.
- ✓Hybridisation (sp, sp2, sp3, sp3d, sp3d2) matches the steric number and fixes the geometry.
- ✓Bond order = (Nb - Na)/2; a higher bond order means a shorter, stronger bond.
- ✓A molecule is polar only if its bond dipoles do not cancel: H2O is polar, but CO2 is non-polar because its two dipoles cancel.
⚡ Quick check
In molecular orbital theory, why is O2 paramagnetic?
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