Chemical Bonding: Shapes and Intermolecular Forces

This guide covers subtopics 3.4 to 3.7 of Topic 3, Chemical bonding, from Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. This is AS Level content.

Before studying this

This resource does not re-teach ionic or covalent bonding from first principles — Ionic, Covalent and Metallic Bonding covers those foundations for IGCSE/O Level, and 9701’s own subtopics 3.1–3.3 (electronegativity, ionic bonding, metallic bonding) build on that before this page begins. What follows is 9701-specific: bonding described in terms of orbitals, not just shared or transferred electron pairs, plus the geometry and intermolecular forces that follow from it.

Syllabus coverage

CAMBRIDGE INTERNATIONAL AS & A LEVEL CHEMISTRY 9701 — AS Level, Topic 3

3.4 Covalent bonding and coordinate (dative covalent) bonding — covalent bonding as electrostatic attraction between nuclei and a shared electron pair, in molecules including H₂, O₂, N₂, Cl₂, HCl, CO₂, NH₃, CH₄, C₂H₆ and C₂H₄; expanded octets in Period 3 elements (SO₂, PCl₅, SF₆); coordinate bonding, including NH₃ + HCl → NH₄⁺ and the Al₂Cl₆ molecule; σ and π bonds from orbital overlap, and how they form in H₂, C₂H₆, C₂H₄, HCN and N₂; hybridisation (sp, sp², sp³); bond energy and bond length, defined and used to compare reactivity.

3.5 Shapes of molecules — using VSEPR theory to state and explain shape and bond angle for BF₃, CO₂, CH₄, NH₃, H₂O, SF₆ and PF₅; predicting the shape and bond angle of analogous molecules and ions.

3.6 Intermolecular forces, electronegativity and bond properties — hydrogen bonding in N–H and O–H containing molecules (ammonia, water); using hydrogen bonding to explain water/ice’s anomalous melting/boiling points, surface tension and solid-vs-liquid density; electronegativity, bond polarity and dipole moments; van der Waals’ forces as a generic term, split into instantaneous dipole–induced dipole (London dispersion) forces and permanent dipole–permanent dipole forces (including hydrogen bonding); the general rule that ionic, covalent and metallic bonding are stronger than intermolecular forces.

3.7 Dot-and-cross diagrams — using dot-and-cross diagrams for ionic, covalent and coordinate bonding, including species from 3.4 and 3.5, and species with an expanded octet or an odd number of electrons.

Orbital overlap: σ and π bonds

A σ (sigma) bond forms by direct, head-on overlap of orbitals between two bonding atoms; a π (pi) bond forms by sideways overlap of adjacent p orbitals, above and below the σ bond. A single bond is always one σ bond alone; a double bond is one σ bond plus one π bond — which is why the second bond in a C=C double bond is weaker and more reactive than the first: sideways overlap is less effective than head-on overlap.

Hybridisation describes how an atom’s orbitals mix before bonding: an sp³ carbon (as in methane or ethane) has four equivalent hybrid orbitals arranged tetrahedrally; an sp² carbon (as in ethene) has three hybrid orbitals in a plane, leaving one unhybridised p orbital free to form a π bond; an sp carbon has two hybrid orbitals arranged linearly, leaving two unhybridised p orbitals for two π bonds.

Shapes of molecules — VSEPR theory

Valence Shell Electron Pair Repulsion (VSEPR) theory says electron pairs around a central atom — bonding and lone pairs alike — arrange themselves as far apart as possible, to minimise repulsion between them. Lone pairs repel more strongly than bonding pairs, which is why replacing a bonding pair with a lone pair compresses the remaining bond angles.

Molecule	Electron pairs (bonding + lone)	Shape	Bond angle
BF₃	3 + 0	Trigonal planar	120°
CO₂	2 + 0	Linear	180°
CH₄	4 + 0	Tetrahedral	109.5°
NH₃	3 + 1	Pyramidal	107°
H₂O	2 + 2	Non-linear (bent)	104.5°
PF₅	5 + 0	Trigonal bipyramidal	120° and 90°
SF₆	6 + 0	Octahedral	90°

The trend from CH₄ (109.5°) to NH₃ (107°) to H₂O (104.5°) is a direct consequence of lone-pair repulsion: each lone pair added in place of a bonding pair squeezes the remaining bond angles a little further.

Intermolecular forces

Every molecule experiences van der Waals’ forces — the generic term for intermolecular forces other than the bonds holding a molecule together. Two kinds matter at AS:

Instantaneous dipole–induced dipole (id-id) forces, also called London dispersion forces, exist between all molecules, arise from momentary, randomly fluctuating electron distributions, and are the only intermolecular force possible between non-polar molecules.

Permanent dipole–permanent dipole (pd-pd) forces exist between polar molecules with a permanent uneven charge distribution. Hydrogen bonding is the strongest kind of permanent dipole–permanent dipole force — it occurs specifically where hydrogen is bonded directly to a small, highly electronegative atom (limited, at this level, to N–H and O–H).

Why water is anomalous

Hydrogen bonding between water molecules explains three properties that would otherwise be surprising for so small a molecule: its relatively high melting and boiling points (extra energy is needed to break the hydrogen bonds, not just weaker van der Waals’ forces, between molecules); its relatively high surface tension (surface molecules are pulled inward by hydrogen bonds with fewer neighbours to balance the pull); and ice being less dense than liquid water (in ice, hydrogen bonds lock molecules into a fixed, open lattice with more empty space than the liquid has, so the same mass occupies a larger volume).

All of this is still weaker than a covalent, ionic or metallic bond — hydrogen bonds and van der Waals’ forces act between molecules, not within them, which is the general rule 3.6 closes on.

Common mistakes

• Calling all intermolecular forces “van der Waals’ forces” as if that meant only the weak dispersion type. The syllabus uses van der Waals’ forces as the umbrella term that includes hydrogen bonding, not something separate from it.
• Predicting a shape without counting lone pairs. Electron-pair count, not just the number of bonded atoms, determines the shape.
• Confusing bond angle compression with bond strength. NH₃’s smaller angle than CH₄ is about lone-pair repulsion, not about the N–H bond being weaker.
• Explaining ice’s lower density using “hydrogen bonds are strong” alone. The density effect specifically needs the open, ordered lattice argument, not just bond strength.
• Drawing a π bond as if it were a second, identical σ bond. They form by a different kind of orbital overlap and behave differently, particularly in reactivity.

Quick revision checklist

• σ bonds (head-on overlap) vs π bonds (sideways overlap); hybridisation sp, sp², sp³
• Coordinate bonding, and the ammonium ion / Al₂Cl₆ examples
• The seven VSEPR shapes and bond angles: BF₃, CO₂, CH₄, NH₃, H₂O, PF₅, SF₆
• Van der Waals’ forces as the umbrella term; id-id (London) vs pd-pd forces
• Hydrogen bonding as a special case of pd-pd forces, limited to N–H/O–H
• Water/ice’s three anomalous properties, each explained by hydrogen bonding
• Dot-and-cross diagrams, including expanded-octet and odd-electron species

Related resources

• Ionic, Covalent and Metallic Bonding — the IGCSE/O Level foundation this resource builds on
• Atomic Structure: Orbitals and Ionisation Energy — orbital theory behind hybridisation
• Chemical Energetics: Hess’s Law and Enthalpy Cycles — bond energy used quantitatively
• Cambridge AS & A Level Chemistry hub

Written against Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. Always check the current syllabus for your examination year.