Bonding and Intermolecular Forces

Ionic bonding happens when electrons jump from one atom to another, creating charged particles that attract each other like magnets. These ionic bonds create massive structures called giant ionic lattices - think of them as 3D networks of alternating positive and negative ions.

Here's what makes ionic compounds special: they have sky-high melting and boiling points because breaking those electrostatic forces requires serious energy. Most dissolve brilliantly in water because the water molecules can pull the ions apart more easily than the ions can hold onto each other.

Covalent bonding works differently - atoms actually share electrons rather than stealing them. The covalent bond forms when electron clouds overlap, creating a strong attraction between the shared electrons and both atomic nuclei. Sometimes one atom provides both electrons in the shared pair, which we call a dative covalent bond.

Quick Tip: Think ionic = transfer, covalent = share!

When drawing ionic equations, remember the golden rules: split anything aqueous (aq) into ions, keep solids/liquids/gases whole, and cancel out identical ions from both sides.

Molecular Shapes and Electronegativity

Molecular geometry isn't random - it follows predictable patterns based on electron pairs. Bonding pairs spread out evenly to minimise repulsion, giving you shapes like linear (180°), trigonal planar (120°), and tetrahedral (109.5°). But lone pairs are bullies - they're more electron-dense and push bonding pairs closer together.

Electronegativity measures how desperately an atom wants to hog electrons in a bond. The Pauling Scale shows this increases as you go up groups and across periods (fluorine is the ultimate electron grabber). Three key factors control this: atomic radius, nuclear charge, and electron shielding.

The electronegativity difference between atoms determines bond type: massive differences (>1.8) create ionic bonds, moderate differences $0.5-1.8$ give polar covalent bonds with partial charges, and tiny differences $0.1-0.5$ produce non-polar covalent bonds.

Remember: Polar dissolves polar - it's chemistry's version of "like attracts like"!

Just because a molecule contains polar bonds doesn't automatically make it polar overall - the shape matters too!

Intermolecular Forces

Intermolecular forces are the weak attractions between separate molecules - much weaker than the bonds within molecules but absolutely crucial for determining physical properties. There are three main types you need to master.

London forces exist between all molecules, even non-polar ones. Electrons constantly move around, creating temporary dipoles that induce dipoles in neighbouring molecules. These forces get stronger with more electrons - that's why noble gases have higher boiling points as you go down the group.

Permanent dipole-dipole forces occur between polar molecules where one end is slightly positive and the other slightly negative. These molecules line up with opposite charges attracting each other.

Hydrogen bonding is the strongest intermolecular force. It happens when hydrogen attached to nitrogen, oxygen, or fluorine gets attracted to lone pairs on these highly electronegative atoms in other molecules.

Key Insight: More electrons = stronger London forces = higher boiling points!

Understanding these forces explains why water has such unusual properties and why some substances mix whilst others don't.