Intermolecular forces, electronegativity and bond properties

Polar bond. When two bonded atoms have different electronegativities, the bonding electrons are pulled towards the more electronegative atom, giving it a partial negative charge ($\delta-$) and the other a partial positive ($\delta+$). This separation of charge is a dipole.

Polar vs non-polar molecules. A molecule with polar bonds is only polar overall if the bond dipoles do not cancel:

• CO₂ has two polar C=O bonds, but they point in opposite directions (linear) → dipoles cancel → non-polar molecule.
• H₂O has two polar O–H bonds, but the molecule is bent → dipoles don't cancel → polar molecule.

So molecular polarity depends on both electronegativity difference and molecular shape (symmetry).

van der Waals' forces is the generic term for all intermolecular forces (forces between molecules, not the bonds within them). There are two types:

1. Instantaneous dipole–induced dipole (id–id) — also called London dispersion forces.

• Random movement of electrons creates a momentary instantaneous dipole in a molecule.
• This induces a dipole in a neighbouring molecule → weak attraction.
• Present in all molecules (including non-polar ones and noble gases).
• Stronger with more electrons (larger molecules/atoms) → higher boiling points. This explains why boiling point rises down the noble gases, down Group 17 (F₂ < Cl₂ < Br₂ < I₂), and up the alkanes.

2. Permanent dipole–permanent dipole (pd–pd).

• Act between molecules that have a permanent dipole (polar molecules).
• Generally stronger than id–id for molecules of similar size.

Hydrogen bonding is a special, stronger case of pd–pd (next section).

Hydrogen bonding occurs when a hydrogen atom is bonded to a highly electronegative atom (N or O — F too in general chemistry), making the H very $\delta+$. This H is attracted to a lone pair on the N or O of a neighbouring molecule.

Requirements: an H bonded to N or O, and a lone pair on an N or O of another molecule. It is a special, strong type of pd–pd force.

Hydrogen bonding explains water's unusual behaviour:

• High melting and boiling points (vs other hydrides like H₂S). Hydrogen bonds are extra forces that need significant energy to overcome, raising the mp/bp.
• High surface tension. Hydrogen bonds at the surface pull molecules together strongly, creating a "skin" that supports small insects and forms droplets.
• Ice is less dense than liquid water (ice floats). In ice, each water molecule forms four hydrogen bonds in an open, regular tetrahedral lattice. This holds the molecules further apart than in liquid water, so solid ice is less dense — water expands on freezing. This is why ice floats and why lakes freeze from the top down.

Putting it all in order (weakest → strongest): $\text{id–id} < \text{pd–pd} < \text{hydrogen bond} \;\lll\; \text{covalent / ionic / metallic bonds}$

Ionic, covalent and metallic bonding are generally much stronger than intermolecular forces. That is why, when a simple molecular substance melts or boils, it is the weak intermolecular forces that break — not the strong covalent bonds inside the molecules. (e.g. boiling water breaks hydrogen bonds, not O–H bonds.)