Covalent Bonding & Shapes

Picture two people, each holding out one hand. Alone, each hand is empty — incomplete. The moment they clasp hands, that single pair of hands belongs to both of them at once. Atoms do exactly this with electrons.

Atoms bond to reach a stable, lower-energy arrangement — usually an octet of 8 valence electrons (a duplet of 2 for H, He). They can transfer electrons (ionic) or, like the clasped hands, share them (covalent).

• Octet rule — atoms gain, lose or share electrons to reach 8 in the valence shell (2 for H).
• Covalent bond — a shared pair of electrons that counts for both atoms, just as a handshake belongs to both people.

If sharing one hand is a single bond, then gripping with two hands is a double bond, and a full three-hand clasp is a triple bond. More hands locked together → a tighter, stronger, shorter hold.

Valence Bond Theory (VBT) explains the grip with orbital overlap. A head-on, axial overlap is a strong sigma (σ) bond — the first handshake. Any extra pairs come from sideways overlap of parallel p-orbitals, the weaker pi (π) bonds.

Usually each partner brings one hand to a handshake. But sometimes a generous host offers both hands while the guest arrives empty-handed. The handshake still forms — but every electron in it came from one atom alone.

• Coordinate (dative) bond — a covalent bond in which both shared electrons come from the same donor atom; the acceptor supplies only an empty orbital.
• Ammonium ion — NH₃ (lone pair on N) + H⁺ → NH₄⁺; nitrogen donates the pair.
• Hydronium ion — H₂O + H⁺ → H₃O⁺; oxygen donates the pair.

Think of the central atom as a wheel hub and its bonds as spokes pushed out as evenly as possible. Hybridisation is the mixing of atomic orbitals of similar energy into an equal number of new, identical hybrid orbitals — the evenly spaced spokes.

Tie a few balloons together at their stems and let go: they push each other as far apart as they can. Electron pairs do exactly the same around a central atom. That is VSEPR — Valence-Shell Electron-Pair Repulsion.

Two balloons spring to opposite sides (linear, 180°); three spread into a flat triangle (120°); four bulge into a tetrahedron (109.5°). A lone-pair balloon is fatter and pushes hardest, so it squeezes the bond angles down.

Go back to the handshakes. A three-hand clasp pulls the two people closest together and is the hardest to break; a one-hand hold is loosest and easiest to part. Bonds behave the same way.

• Bond length — the distance between two bonded nuclei. Triple < double < single.
• Bond energy — the energy needed to break the bond. Triple > double > single.
• Bond angle — the angle between two bonds at the central atom.

Two equal teams on a tug-of-war rope keep the flag in the middle — a non-polar bond between identical atoms (H₂, Cl₂). But if one team is stronger (more electronegative), it drags the flag toward itself: the shared electrons sit closer to that atom, giving it a partial negative charge (δ−) and the other δ+. That lopsided pull is a dipole.

Shape decides molecular polarity. In linear CO₂ the two equal pulls point opposite ways and cancel — non-polar. In bent water the two pulls point the same general way and add up, leaving water a tiny magnet-like molecule with a positive and a negative end.

Molecular Orbital Theory (MOT) treats the whole molecule: atomic orbitals combine into bonding MOs (lower energy, they hold the molecule together) and antibonding MOs (higher energy, they pull it apart).

1. Octet rule; ionic vs covalent vs coordinate (dative) bonds.
2. Single / double / triple bonds; σ (head-on) vs π (sideways) overlap.
3. Hybridisation sp³ / sp² / sp — the evenly spread spokes.
4. VSEPR — balloons repel into the five shapes; lone pairs shrink the angle.
5. Bond parameters — length, energy, angle; polarity & dipole; MOT & bond order.