Atoms bond to reach a lower-energy arrangement — usually a full outer shell. Class 11 asks four things: which kind of bond forms, what shape the molecule takes, what orbitals are involved, and why some molecules are paramagnetic. Each has a rule you can apply in under a minute.
Kinds of bond
- Ionic — electrons are transferred from a metal to a non-metal; the ions attract. Example: NaCl. Favoured by low ionisation energy (metal), high electron affinity (non-metal) and high lattice energy.
- Covalent — electrons are shared between non-metals. Example: Cl₂, H₂O, CH₄. Single, double (O₂) or triple (N₂) bonds.
- Coordinate (dative) — both shared electrons come from one atom. Examples: NH₄⁺ (N gives its lone pair to H⁺), H₃O⁺, the N→O in HNO₃.
- Polar covalent — shared unequally because of electronegativity difference; H–Cl has a dipole. Whether the molecule is polar depends on shape: CO₂ is linear and non-polar despite polar bonds; H₂O is bent and polar.
Fajan's rules predict covalent character in an "ionic" compound: a small, highly charged cation and a large anion polarise strongly. So LiI is more covalent than LiF; AlCl₃ is more covalent than NaCl.
Lewis structures and formal charge
Count valence electrons, place the least electronegative atom at the centre (never H), connect with single bonds, distribute the remaining electrons to satisfy octets, and make multiple bonds if the centre is short.
Formal charge = valence electrons − non-bonding electrons − ½(bonding electrons). The best structure has formal charges closest to zero. In CO₂ the structure O=C=O has all zeros; O≡C–O does not.
Octet exceptions: BF₃ (6 electrons on B), PCl₅ and SF₆ (expanded octet), NO and NO₂ (odd electron).
VSEPR — predicting shape
Electron pairs around the central atom repel and spread out as far as possible. Count bond pairs + lone pairs; lone pairs repel more, so they squeeze the bond angles.
| Domains | Lone pairs | Shape | Example | Angle |
|---|---|---|---|---|
| 2 | 0 | Linear | BeCl₂, CO₂ | 180° |
| 3 | 0 | Trigonal planar | BF₃ | 120° |
| 3 | 1 | Bent | SO₂ | <120° |
| 4 | 0 | Tetrahedral | CH₄ | 109.5° |
| 4 | 1 | Trigonal pyramidal | NH₃ | 107° |
| 4 | 2 | Bent | H₂O | 104.5° |
| 5 | 0 | Trigonal bipyramidal | PCl₅ | 120°, 90° |
| 5 | 1 | See-saw | SF₄ | |
| 5 | 2 | T-shaped | ClF₃ | |
| 5 | 3 | Linear | XeF₂ | 180° |
| 6 | 0 | Octahedral | SF₆ | 90° |
| 6 | 1 | Square pyramidal | BrF₅ | |
| 6 | 2 | Square planar | XeF₄ | 90° |
Repulsion order: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair. That is why the angle drops from 109.5° (CH₄) to 107° (NH₃) to 104.5° (H₂O).
Hybridisation — which orbitals
The number of electron domains gives the hybridisation directly:
- 2 → sp (linear): BeCl₂, C₂H₂, CO₂.
- 3 → sp² (trigonal planar): BF₃, C₂H₄, SO₂ (with a lone pair).
- 4 → sp³ (tetrahedral): CH₄, NH₃, H₂O.
- 5 → sp³d (trigonal bipyramidal): PCl₅.
- 6 → sp³d² (octahedral): SF₆.
Bonds: a single bond is one σ; a double bond is σ + π; a triple bond is σ + 2π. In ethene each carbon is sp²; the C=C has one σ (sp²–sp²) and one π (p–p). In ethyne each carbon is sp with two π bonds.
Shortcut for the steric number: (valence electrons of central atom + number of monovalent atoms attached ± charge)/2. For NH₃: (5 + 3)/2 = 4 → sp³. For NH₄⁺: (5 + 4 − 1)/2 = 4 → sp³.
Molecular orbital theory — why O₂ is paramagnetic
Atomic orbitals combine into bonding (lower energy) and antibonding (higher, marked *) molecular orbitals. Fill them in order, and:
Bond order = (bonding electrons − antibonding electrons)/2.
Order for O₂, F₂: σ1s, σ1s, σ2s, σ2s, σ2p_z, π2p_x = π2p_y, π2p_x = π2p_y, σ*2p_z. Order for B₂, C₂, N₂ (and lighter): the π2p pair lies below σ2p_z.
- N₂: 14 electrons, bond order (10 − 4)/2 = 3; diamagnetic.
- O₂: 16 electrons, bond order (10 − 6)/2 = 2; the last two electrons go singly into the two π* orbitals — two unpaired electrons, paramagnetic.
- O₂⁺: bond order 2.5 (stronger, shorter than O₂); O₂⁻: 1.5; O₂²⁻: 1.
- He₂: bond order 0 — does not exist.
Hydrogen bonding
An H attached to N, O or F attracts a lone pair on another N, O or F. It explains why water boils at 100 °C while H₂S is a gas, why ice floats (open lattice), and why o-nitrophenol is more volatile than p-nitrophenol (intramolecular vs intermolecular H-bonding).
How it is tested
- Shape and hybridisation of a given molecule (NH₃, SF₄, XeF₄ are favourites).
- Bond order and magnetism from MOT (O₂ species, N₂ vs N₂⁺).
- Compare bond angles (CH₄ > NH₃ > H₂O) with the lone-pair reason.
- Dipole moment: why CO₂ and BF₃ are zero, why NH₃ is larger than NF₃.
- Fajan's rules: which of two compounds is more covalent.