1. Primary Chemical Bonding
Chemical bonding arises from electrostatic attractions between charged particles, driving atoms towards lower-energy configurations.
| Bonding Type | Definition & Nature of Force | Typical Examples | Macroscopic Properties |
|---|---|---|---|
| Ionic | Electrostatic attraction between oppositely charged ions in a giant 3D lattice. | \(\text{NaCl}, \text{MgO}, \text{CaF}_2\) | High melting points, brittle, conduct electricity when molten or dissolved in water. |
| Covalent | Shared pair of electrons between non-metal atoms. | \(\text{Cl}_2, \text{H}_2\text{O}, \text{CH}_4\) | Low melting points for simple molecules; very high for giant covalent macromolecular structures. |
| Coordinate (Dative) | Shared pair of electrons where both electrons are donated by the same atom. | \(\text{NH}_4^+, \text{H}_3\text{O}^+, \text{Al}_2\text{Cl}_6\) | Identical in length, strength, and chemical behavior to ordinary covalent bonds once formed. |
| Metallic | Electrostatic attraction between positive metal cations and a delocalised sea of electrons. | \(\text{Na}, \text{Mg}, \text{Fe}, \text{Cu}\) | High thermal and electrical conductivity, malleable and ductile due to layers of ions sliding. |
2. Electronegativity and Bond Polarity
The power of an atom to attract the electron density in a covalent bond towards itself (Pauling scale).
Electronegativity increases across a period (greater nuclear charge, smaller atomic radius) and decreases down a group (increased electron shielding and distance). Fluorine is the most electronegative element (4.0 on Pauling scale).
Bond Polarity vs Molecular Polarity
A bond between atoms with differing electronegativities is polar (\(\text{C}^{\delta+} - \text{O}^{\delta-}\)). However, whether the whole molecule is polar depends on molecular symmetry:
- Carbon dioxide (\(\text{CO}_2\)): Linear shape (\(\text{O}=\text{C}=\text{O}\)). The two equal bond dipoles act in opposite directions and cancel out completely. \(\text{CO}_2\) is non-polar.
- Water (\(\text{H}_2\text{O}\)): Bent shape due to two lone pairs. The bond dipoles do not cancel, giving a permanent molecular dipole moment.
3. Intermolecular Forces
Intermolecular forces exist between independent covalent molecules, dictating physical properties such as boiling point and viscosity:
- London Dispersion Forces (Induced Dipole-Dipole): Present between all molecules. Constant electron movement creates an instantaneous temporary dipole, which induces an opposing dipole in an adjacent molecule. Strength increases with greater electron count and larger molecular surface area.
- Permanent Dipole-Dipole Forces: Electrostatic attractions between the permanent \(\delta+\) end of one polar molecule and the \(\delta-\) end of another. Stronger than London dispersion forces for molecules of comparable electron count.
- Hydrogen Bonding: The strongest intermolecular force. Requires a hydrogen atom covalently bonded to a highly electronegative atom (\(\text{F}, \text{O}, \text{N}\)) with at least one lone pair of electrons.
- Anomalously high boiling points: \(\text{H}_2\text{O}\), \(\text{HF}\), and \(\text{NH}_3\) boil at much higher temperatures than Group 16, 17, and 15 hydrides because significant energy is required to break extensive hydrogen bonds.
- Density of Ice: In ice, each water molecule forms four hydrogen bonds in an open hexagonal lattice. Molecules are held further apart than in liquid water, making ice less dense than liquid water (ice floats).
4. VSEPR Molecular Geometry
Valence Shell Electron Pair Repulsion theory states that electron pairs around a central atom repel each other and adopt positions of maximum separation to minimise electrostatic repulsion.
| Total Pairs | Bonding Pairs | Lone Pairs | Shape Name | Bond Angle | Example |
|---|---|---|---|---|---|
| 2 | 2 | 0 | Linear | 180° | \(\text{BeCl}_2, \text{CO}_2\) |
| 3 | 3 | 0 | Trigonal planar | 120° | \(\text{BF}_3, \text{AlCl}_3\) |
| 3 | 2 | 1 | Bent (V-shaped) | 118° | \(\text{SO}_2\) |
| 4 | 4 | 0 | Tetrahedral | 109.5° | \(\text{CH}_4, \text{NH}_4^+\) |
| 4 | 3 | 1 | Trigonal pyramidal | 107° | \(\text{NH}_3, \text{H}_3\text{O}^+\) |
| 4 | 2 | 2 | Bent | 104.5° | \(\text{H}_2\text{O}, \text{H}_2\text{S}\) |
| 5 | 5 | 0 | Trigonal bipyramidal | 90° & 120° | \(\text{PCl}_5\) |
| 5 | 4 | 1 | Seesaw | 87° & 102° | \(\text{SF}_4\) |
| 5 | 3 | 2 | T-shaped | 88° | \(\text{ClF}_3\) |
| 6 | 6 | 0 | Octahedral | 90° | \(\text{SF}_6\) |
| 6 | 4 | 2 | Square planar | 90° | \(\text{XeF}_4\) |
5. Crystal Structures
Diamond
Giant macromolecular lattice where each carbon atom forms four strong covalent bonds in a tetrahedral arrangement (\(109.5^\circ\)). Extremely hard, very high melting point (\(>3800\text{ K}\)), electrical insulator.
Graphite
Each carbon forms three covalent bonds in planar hexagonal layers (\(120^\circ\)). One delocalised electron per carbon travels freely between layers, making graphite an electrical conductor. Weak London forces between sheets allow layers to slide easily (solid lubricant).
6. Practice Questions
Step 1: Count valence electron pairs around central Xenon:
- Xe is in Group 0 (8 valence electrons).
- 4 single bonds to Fluorine atoms contribute 4 electrons: \(8 + 4 = 12\) electrons = 6 electron pairs.
- 4 bonding pairs and 2 lone pairs.
Step 2: Apply VSEPR reasoning:
- 6 electron pairs arrange octahedrally to minimise repulsion.
- The 2 lone pairs occupy opposite axial positions (\(180^\circ\) apart) to minimise intense lone pair-lone pair repulsion.
- The resulting molecular shape is square planar.
- Bond angle: exactly 90 degrees.