Molecular Shapes, Bond Polarity and Intermolecular Forces

This guide covers sections 3.1.3.5 to 3.1.3.7, the shapes, polarity and intermolecular-forces half of the Bonding section (3.1.3) in OxfordAQA International AS and A-level Chemistry 9620. It follows Ionic, covalent, metallic bonding and the four crystal types and moves from how atoms bond to the three-dimensional shapes and polarity that bonding produces.

Shapes of simple molecules and ions

You need to understand molecular shape through electron-pair repulsion theory. Bonding pairs and lone (non-bonding) pairs of electrons behave as charge clouds that repel one another, and the electron pairs in the outer shell of a central atom arrange themselves as far apart as possible to minimise this repulsion – this single principle is what determines a molecule’s overall shape.

A specific hierarchy of repulsion strength applies: lone pair-lone pair repulsion is greater than lone pair-bond pair repulsion, which is greater than bond pair-bond pair repulsion. This hierarchy explains why bond angles shrink slightly whenever a lone pair is present – for example, why ammonia’s H-N-H bond angle (107°) is smaller than methane’s perfectly tetrahedral 109.5°, despite both having four electron pairs around the central atom. You need to be able to name and explain the shapes of, and bond angles in, simple molecules and ions with up to six electron pairs (including lone pairs) surrounding the central atom – covering the standard shape sequence from linear (2 pairs) through trigonal planar (3), tetrahedral (4), trigonal bipyramidal (5), to octahedral (6), plus their lone-pair-adjusted variants such as bent, pyramidal and see-saw shapes.

Bond polarity

You need to understand electronegativity as the power of an atom to attract the shared pair of electrons in a covalent bond. When two bonded atoms differ in electronegativity, the electron distribution in that bond becomes unsymmetrical, producing a polar covalent bond, which may in turn give the whole molecule a permanent dipole. You need to be able to use partial charges (δ+ and δ−) to show that a bond is polar, and – a frequently tested subtlety – explain why some molecules containing polar bonds do not have an overall permanent dipole. This happens when a molecule’s shape is symmetrical enough that the individual bond dipoles cancel out vectorially, as in carbon dioxide (linear, two identical polar C=O bonds pointing in opposite directions) or tetrachloromethane (tetrahedral, four identical polar C-Cl bonds arranged symmetrically).

Forces between molecules

You need to know three types of intermolecular force: permanent dipole-dipole forces, induced dipole-dipole forces (also called van der Waals, dispersion, or London forces), and hydrogen bonding. You need to be able to explain the existence of these forces between both familiar and unfamiliar molecules, and explain how a substance’s melting and boiling points are influenced by the relative strength of whichever forces are present.

Two specific applications are named directly in the specification: the importance of hydrogen bonding in explaining the low density of ice (relative to liquid water, an unusual property most substances do not show), and the anomalous boiling points of compounds that can hydrogen bond compared with similarly sized molecules that cannot.

How the three forces compare in strength

Induced dipole-dipole (dispersion) forces exist between all molecules, regardless of polarity, and generally increase in strength with increasing number of electrons or surface area of contact. Permanent dipole-dipole forces add to this baseline in polar molecules. Hydrogen bonding is the strongest of the three intermolecular forces (though still much weaker than a covalent or ionic bond), and requires a very specific arrangement: a hydrogen atom bonded directly to a small, highly electronegative atom (fluorine, oxygen or nitrogen) that also carries a lone pair, allowing that hydrogen to form an additional weak bond to a lone pair on a neighbouring molecule.

A worked example

Water and hydrogen sulfide are structurally similar molecules (both have two hydrogen atoms bonded to a central atom from Group 6), yet water boils at 100°C while hydrogen sulfide boils at −60°C. The explanation lies entirely in intermolecular forces: oxygen is small and highly electronegative, so O-H bonds allow water molecules to hydrogen bond to each other, while sulfur is larger and less electronegative, so H₂S molecules can only interact through weaker dispersion and permanent dipole-dipole forces. Overcoming water’s additional hydrogen bonding on boiling requires substantially more energy, producing the large boiling-point gap between two molecules of similar size and shape.

How to approach it

When predicting a molecular shape, count electron pairs (bonding plus lone) around the central atom first, identify the base shape from that count, then adjust the bond angle downward for each lone pair present using the repulsion hierarchy – working in this fixed order avoids the common error of guessing a shape from memory without justifying it from electron-pair count. For polarity questions, always check molecular symmetry before concluding a molecule is non-polar just because it contains polar bonds – the correct answer requires both the presence of polar bonds and an explanation of why their dipoles do, or do not, cancel by symmetry. For intermolecular force questions, identify all forces present systematically: dispersion forces are always present, then check for permanent dipoles from bond polarity and molecular shape, then check specifically for the H-F, H-O or H-N arrangement required for hydrogen bonding – a molecule can have more than one type of intermolecular force acting simultaneously, and exam answers that name only the strongest force while ignoring the others present are frequently incomplete.

Official syllabus

OxfordAQA, International AS and A-level Chemistry (9620) specification, Version 4.3, for International AS and A-level exams May/June 2020 onwards: official specification PDF, sections 3.1.3.5-3.1.3.7 “Bonding”. Verified 2026-09-02.