Types of bonding revisited
| Bond | Formed by | Example |
|---|---|---|
| Ionic | Transfer of electrons; electrostatic attraction between ions in a lattice | NaCl, MgO |
| Covalent | Sharing of electron pairs | H₂O, CH₄ |
| Coordinate (dative covalent) | Both electrons of the shared pair come from one atom | NH₄⁺, H₃O⁺ |
| Metallic | Positive ions in a sea of delocalised electrons | Cu, Fe |
Ionic bonding is strongest when the ions are small and highly charged (large lattice energy). When a small positive ion has a large charge it can distort (polarise) the electron cloud of a large negative ion, giving the bond some covalent character (Fajans’ rules).
Shapes of molecules: VSEPR
Pairs of electrons around a central atom repel one another and move as far apart as possible. The valence shell electron pair repulsion (VSEPR) theory predicts the shape from the number of bonding pairs and lone pairs. Lone pairs repel more strongly than bonding pairs, so they squeeze the bond angles.
| Electron pairs | Lone pairs | Shape | Bond angle | Example |
|---|---|---|---|---|
| 2 | 0 | Linear | 180° | BeCl₂, CO₂ |
| 3 | 0 | Trigonal planar | 120° | BF₃ |
| 4 | 0 | Tetrahedral | 109.5° | CH₄ |
| 4 | 1 | Pyramidal | 107° | NH₃ |
| 4 | 2 | Bent (V-shaped) | 104.5° | H₂O |
| 5 | 0 | Trigonal bipyramidal | 90° and 120° | PCl₅ |
| 6 | 0 | Octahedral | 90° | SF₆ |
Hybridisation
To explain the shape of methane, an atom of carbon mixes (hybridises) its 2s and 2p orbitals to make four equal sp³ orbitals pointing to the corners of a tetrahedron. In ethene each carbon has three sp² orbitals in a plane (120°) and one p orbital left over; the sideways overlap of two p orbitals forms the π bond of the double bond. In ethyne each carbon is sp hybridised (180°) and has two π bonds. A σ bond is formed by end-on overlap; a π bond by sideways overlap.
| Hybridisation | Shape | Example | Bonds |
|---|---|---|---|
| sp³ | Tetrahedral, 109.5° | CH₄, C₂H₆ | All σ |
| sp² | Trigonal planar, 120° | C₂H₄, benzene | σ and one π |
| sp | Linear, 180° | C₂H₂, CO₂ | σ and two π |
Polarity and intermolecular forces
In a bond between atoms of different electronegativity the shared pair is pulled towards the more electronegative atom, giving a polar bond (δ⁺ and δ⁻). A molecule is polar only if its shape does not cancel the bond polarities: water is polar, carbon dioxide is not.
| Force | Where | Relative strength |
|---|---|---|
| London (dispersion) forces | Between all molecules; increase with the number of electrons | Weakest |
| Permanent dipole–dipole | Between polar molecules | Medium |
| Hydrogen bonding | H bonded to N, O or F, attracted to a lone pair on N, O or F in another molecule | Strongest of the three |
| Covalent, ionic, metallic bonds | Within the substance | Much stronger than intermolecular forces |
Choose the electronegativity difference between the two atoms. The percentage ionic character is estimated from Pauling’s formula: 1 − e^(−0.25 × Δχ²).
Key terms
- VSEPR
- A theory that predicts molecular shape from repulsion between electron pairs.
- Hybridisation
- The mixing of atomic orbitals to make new equivalent orbitals.
- Dative covalent bond
- A covalent bond in which both shared electrons come from one atom.
- Hydrogen bond
- A strong dipole–dipole attraction involving H bonded to N, O or F.
Summary
- Ionic, covalent, dative and metallic bonds; polarisation gives ionic bonds covalent character.
- VSEPR: lone pairs repel more than bonding pairs; know the shapes and angles.
- sp³, sp² and sp hybridisation explain σ and π bonds in alkanes, alkenes and alkynes.
- Intermolecular forces: London < dipole–dipole < hydrogen bonding.