Chemical Bonding and Molecular Structure
Chemical Bonding and Molecular Structure
This domain examines why substances have particular structures and properties: ionic lattices, covalent molecules, polar bonds, molecular shapes and intermolecular forces. The exam often asks you to predict a shape and then explain polarity or boiling point.
Core rules
- Ionic bonding involves electrostatic attraction between ions; covalent bonding shares electron pairs. A coordinate bond is a covalent bond where one atom supplies both electrons.
- VSEPR: electron pairs repel and adopt positions that minimise repulsion. Count bonding pairs and lone pairs around the central atom. Lone pairs repel more strongly than bond pairs, reducing bond angles.
- Shapes: 2 electron domains linear; 3 trigonal planar; 4 tetrahedral; 3 bonds + 1 lone pair trigonal pyramidal; 2 bonds + 2 lone pairs bent.
- Intermolecular forces from weakest: induced dipole-dipole, permanent dipole-dipole, hydrogen bonding. Hydrogen bonds form when H is bonded to N, O or F and interacts with a lone pair on N, O or F.
Worked example
Ammonia, NH3, has three N–H bonds and one lone pair on nitrogen. Four electron domains give a tetrahedral arrangement, but the shape is trigonal pyramidal. N–H bonds are polar, and the lone-pair direction means the bond dipoles do not cancel, so NH3 is polar. It forms hydrogen bonds, giving a higher boiling point than PH3.
Traps
A molecule can contain polar bonds and still be non-polar if dipoles cancel, as in CO2. Do not count a double bond as two separate electron domains in VSEPR. When explaining conductivity or melting point, link the structure to whether ions or electrons are free to move.