Why atoms share electrons, how bonds form (VBT & hybridisation), and how VSEPR predicts the shape of a molecule — exam-focused for the BIEK / Sindh Board paper. Read it straight through, or open the interactive lecture to watch bonds and shapes form.
1 — Why atoms combine: the octet rule
Atoms bond to reach a stable, lower-energy electron arrangement — usually the configuration of the nearest noble gas.
- Octet rule — atoms tend to gain, lose or share electrons so as to have 8 electrons in their valence shell (2 for H, He — the duplet).
- Chemical bond — the attractive force that holds atoms together in a molecule or compound.
| Bond type | How it forms | Example |
|---|
| Ionic (electrovalent) | complete transfer of electrons (metal → non-metal) | NaCl |
| Covalent | sharing of electron pairs (non-metals) | H₂, Cl₂, H₂O |
| Coordinate (dative) | shared pair donated by one atom | NH₄⁺, H₃O⁺ |
2 — The covalent bond
- Covalent bond — a bond formed by the mutual sharing of one or more electron pairs between two atoms, so each attains a noble-gas configuration.
Single, double & triple bonds
| Shared pairs | Bond | Example |
|---|
| 1 pair (2 e⁻) | single (—) | H—H, Cl—Cl |
| 2 pairs (4 e⁻) | double (=) | O=O, CH₂=CH₂ |
| 3 pairs (6 e⁻) | triple (≡) | N≡N, HC≡CH |
Lewis (dot-and-cross) structures show the shared and lone pairs. e.g. water: O with two bond pairs (to H) and two lone pairs.
3 — Coordinate (dative) covalent bond
A covalent bond in which both shared electrons come from the same atom (the donor); the acceptor supplies an empty orbital. Once formed, it is identical to an ordinary covalent bond.
- Ammonium ion — NH₃ (lone pair on N) + H⁺ → NH₄⁺; N donates the pair.
- Hydronium ion — H₂O + H⁺ → H₃O⁺; O donates the pair.
Shown with an arrow (→) pointing from donor to acceptor.
4 — Valence Bond Theory (VBT)
VBT (Heitler–London, Pauling) says a covalent bond forms by the overlap of half-filled atomic orbitals. Greater overlap → stronger bond. Two kinds of overlap give two kinds of bond:
- Sigma (σ) bond — head-on (axial) overlap (s–s, s–p, or p–p end-to-end). Strong; every single bond is a σ bond.
- Pi (π) bond — sideways (lateral) overlap of parallel p-orbitals. Weaker; found in double/triple bonds in addition to one σ.
Counting bondssingle = 1σ · double = 1σ + 1π · triple = 1σ + 2π
5 — Hybridisation
- Hybridisation — the intermixing of atomic orbitals of similar energy to form an equal number of new, identical hybrid orbitals oriented for maximum overlap.
| Hybridisation | Orbitals mixed | Geometry · angle | Example |
|---|
| sp³ | 1 s + 3 p | tetrahedral · 109.5° | CH₄, NH₃, H₂O |
| sp² | 1 s + 2 p | trigonal planar · 120° | BF₃, C₂H₄ |
| sp | 1 s + 1 p | linear · 180° | BeCl₂, C₂H₂ |
Methane (CH₄): carbon's 2s and three 2p mix into four sp³ orbitals pointing to the corners of a tetrahedron (109.5°).
6 — VSEPR theory
Valence-Shell Electron-Pair Repulsion theory predicts molecular shape: the electron pairs around the central atom arrange themselves to be as far apart as possible, minimising repulsion.
Repulsion orderlone pair–lone pair > lone pair–bond pair > bond pair–bond pair
Lone pairs repel more strongly, so they squeeze the bond angle down (e.g. CH₄ 109.5° → NH₃ 107° → H₂O 104.5°).
7 — Shapes of molecules
| Bond pairs | Lone pairs | Shape | Angle | Example |
|---|
| 2 | 0 | linear | 180° | BeCl₂, CO₂ |
| 3 | 0 | trigonal planar | 120° | BF₃ |
| 4 | 0 | tetrahedral | 109.5° | CH₄ |
| 3 | 1 | trigonal pyramidal | 107° | NH₃ |
| 2 | 2 | bent (angular) | 104.5° | H₂O |
8 — Bond parameters
- Bond length — the average distance between the nuclei of two bonded atoms. Triple < double < single.
- Bond energy — the energy required to break one mole of a particular bond in the gaseous state. Triple > double > single.
- Bond angle — the angle between two bonds at the central atom.
| Bond | Length (pm) | Energy (kJ/mol) |
|---|
| C—C | 154 | 347 |
| C=C | 134 | 614 |
| C≡C | 120 | 839 |
9 — Polar & non-polar covalent bonds
- Non-polar covalent bond — equal sharing between identical atoms (e.g. H₂, Cl₂); zero electronegativity difference.
- Polar covalent bond — unequal sharing; the more electronegative atom pulls the pair, gaining a partial negative charge (δ−), the other δ+ (e.g. H—Cl).
Dipole momentμ = q × d (unit: debye, D)
Shape decides molecular polarity: CO₂ has polar C=O bonds but is linear → the dipoles cancel → non-polar. H₂O has polar O—H bonds and is bent → dipoles add → polar.
10 — Molecular Orbital Theory (MOT) — an overview
MOT treats the whole molecule: atomic orbitals combine to form molecular orbitals spread over both atoms.
- Bonding MO (σ, π) — lower energy; electrons here stabilise the molecule.
- Antibonding MO (σ*, π*) — higher energy; electrons here weaken the bond.
Bond orderB.O. = ½ ( bonding e⁻ − antibonding e⁻ )
Success: MOT explains why O₂ is paramagnetic (two unpaired electrons) — VBT cannot.
11 — Ionic vs covalent: a comparison
| Property | Ionic compounds | Covalent compounds |
|---|
| Particles | ions in a lattice | molecules |
| Melting / boiling point | high | usually low |
| Electrical conductivity | conduct when molten / aqueous | generally non-conductors |
| Solubility | soluble in water (polar) | soluble in organic (non-polar) solvents |
12 — Worked & reasoning questions
predict the shape
Predict the shape and bond angle of NH₃.
N has 5 valence e⁻ → 3 bond pairs (to H) + 1 lone pair = 4 pairs (sp³)
4 pairs → tetrahedral arrangement, but 1 lone pair → trigonal pyramidal, ~107°
bond order
Bond order of N₂ (bonding 10, antibonding 4)?
B.O. = ½(10 − 4) = 3 (a triple bond)
polarity
Why is CO₂ non-polar but H₂O polar?
Both have polar bonds; CO₂ is linear so the two dipoles cancel; H₂O is bent so they add → CO₂ non-polar, H₂O polar
13 — Exam recap
- Octet rule; ionic vs covalent vs coordinate bonds.
- Single/double/triple bonds; Lewis structures.
- VBT: σ (head-on) vs π (sideways) overlap; bond counting.
- Hybridisation sp³/sp²/sp with geometry and angle.
- VSEPR; the five common shapes and how lone pairs shrink the angle.
- Bond parameters; polarity & dipole moment; MOT bond order.