3.6 Intermolecular forces and bond properties

Hydrogen bonding is a strong intermolecular attraction and a special case of permanent dipole–permanent dipole attraction. For the assessed N–H and O–H examples, the donor molecule has a highly polar bond to hydrogen and the neighbouring molecule supplies a lone pair on an electronegative nitrogen or oxygen atom.

The electronegative atom pulls bonding electrons away from hydrogen, giving H a substantial δ+ charge. That hydrogen is attracted to a lone pair on N or O in another molecule; the attraction is strongest when the covalent and hydrogen bonds are close to linear. Ammonia and water therefore form intermolecular hydrogen-bond networks, with the number of possible links limited by available N–H/O–H hydrogens and lone pairs.

In water, hydrogen bonds hold molecules together, so more energy is needed to separate them: melting and boiling points are relatively high. At the surface, neighbouring molecules pull surface molecules inward, producing high surface tension. In ice, a rigid, more open hydrogen-bonded arrangement leaves molecules farther apart than in liquid water, so ice has a lower density.

Do not call the O–H or N–H covalent bond itself a hydrogen bond: the hydrogen bond is the intermolecular attraction to a neighbouring lone pair. Do not infer hydrogen bonding from the presence of hydrogen alone, and do not replace the network/packing explanation with the claim that all solids are denser than their liquids.

Electronegativity is an atom's tendency to attract the shared electron pair in a covalent bond. Equal or nearly equal attraction gives a non-polar bond; unequal attraction shifts electron density towards the more electronegative atom, leaving that end δ− and the other end δ+.

A bond dipole records the direction of this charge separation, with the dipole arrow pointing towards the partially negative end. A larger electronegativity difference gives a more polar bond, but bond polarity is a property of the individual bond before the whole molecular arrangement is considered.

To decide whether a molecule has an overall dipole, identify the polarity and direction of every polar bond, then use the molecular shape to see whether the dipoles cancel. In CH₃Cl they do not cancel, so the molecule is polar; in the symmetric CCl₄ arrangement they cancel, so the molecule is non-polar despite its polar C–Cl bonds.

Do not infer overall molecular polarity from one bond or from the presence of polar bonds alone. Keep δ+ and δ− on the correct atoms, and distinguish a polar bond from a polar molecule whose net dipole depends on three-dimensional arrangement.

Van der Waals’ forces are intermolecular attractions between molecular entities. The two assessed types are instantaneous dipole–induced dipole (London dispersion) forces and permanent dipole–permanent dipole forces; hydrogen bonding is a special permanent-dipole case and is treated separately for the detailed water consequences.

In every atom or molecule, electron motion can momentarily make the charge cloud uneven. The temporary dipole then induces an opposite dipole in a neighbour, so δ+ and δ− ends attract. These forces generally become stronger as electron number increases and as molecules have more close-contact positions.

A polar molecule has a lasting charge separation, so neighbouring molecules can align opposite partial charges and form permanent dipole–permanent dipole attractions. For small molecules with the same electron count, this interaction is stronger than London dispersion alone; for example, propanone has stronger intermolecular attraction and a higher boiling point than non-polar butane.

Keep the scale clear: covalent bonds are intramolecular, whereas van der Waals’ forces act between molecules. Do not call every intermolecular force a hydrogen bond, treat an instantaneous dipole as permanent, or compare force types without considering the stated electron count and molecular contact.

Intramolecular forces hold a particle together; intermolecular forces act between neighbouring molecules. Ionic bonding is attraction between oppositely charged ions in a lattice, covalent bonding holds shared electron pairs between atoms, and metallic bonding is attraction between positive metal ions and delocalised electrons.

Intermolecular forces include instantaneous dipole–induced dipole attraction, permanent dipole–permanent dipole attraction and hydrogen bonding as a special permanent-dipole case. These forces change how molecular entities attract one another but do not replace the covalent bonds within each molecule.

In general, the forces holding particles together within a substance are stronger than the attractions between molecules. The source-supported comparison places covalent bonding at the strong end and instantaneous dipole–induced dipole attraction at the weak end; use this as a general boundary, not as a universal numerical ranking for every substance.

When explaining a property or a change of state, first identify whether the relevant attraction is within a particle or between particles. Do not describe an intermolecular force as ionic, covalent or metallic bonding, and do not imply that breaking intermolecular attractions breaks the covalent molecule itself.