Ionic Lattices and Covalent Bonding

Ionic lattices are incredibly tough structures because of the strong electrostatic forces holding positive and negative ions together. This is why they have such high melting and boiling points - you need loads of energy to break these attractions apart.

These compounds become electrical conductors when dissolved in water or melted because the ions can finally move freely. Water is particularly good at dissolving ionic compounds since the positive hydrogen atoms attract negative ions whilst the negative oxygen atoms pull positive ions away from the lattice.

Covalent bonds form when atomic orbitals overlap, creating attraction between the shared electron pair and both nuclei. You'll encounter different types: sigma bonds $formed by s-s, s-p, or head-on p-p overlap$ and pi bonds $formed by sideways p-p overlap$. Single bonds contain one sigma bond, whilst double bonds have one sigma plus one pi bond.

Key Insight: Dative covalent bonds occur when one atom donates both electrons to form the pair - like in NH₄⁺ where ammonia accepts a hydrogen ion.

Molecular Shapes and Bond Angles

Predicting molecular geometry becomes straightforward once you understand electron pair repulsion. Molecules arrange themselves to minimise repulsion between electron pairs, creating predictable shapes with specific bond angles.

Linear molecules like CO₂ have 180° bond angles, whilst trigonal planar shapes (like BF₃) show 120° angles. The common tetrahedral arrangement gives you 109.5° angles, but this changes when lone pairs get involved.

Lone pairs are electron bullies - they push bonding pairs closer together. Pyramidal shapes (like NH₃) start tetrahedral but lone pairs reduce the bond angle to about 107°. Bent molecules like water have even smaller angles (104.5°) because two lone pairs create extra repulsion.

Octahedral molecules have six bonding pairs arranged with 90° bond angles, creating a shape that looks like two square pyramids stuck together.

Exam Tip: Remember that lone pairs repel more strongly than bonding pairs - subtract about 2.5° for each lone pair from the standard tetrahedral angle.