Difficulty: Intermediate | Prerequisites: Bonding fundamentals (Part 1), Lewis structures and formal charge (Part 2)

Big Picture

This topic builds on Lewis structures to explain the three-dimensional shapes molecules adopt, why some molecules are polar even when they contain polar bonds, and how molecules interact with each other without forming bonds. You need to be comfortable drawing Lewis structures and counting electron groups before tackling VSEPR geometry. The intermolecular forces section explains physical properties like boiling points and solubility, which come up constantly in later chemistry and biology courses.

TL;DR

VSEPR theory predicts molecular shapes based on electron-pair repulsion around the central atom. Molecular geometry (the shape you see) can differ from electron geometry (which includes lone pairs). Sigma and pi bonds differ in orbital overlap and rotational freedom. Intermolecular forces, from weakest to strongest: London dispersion, dipole-dipole, hydrogen bonding, each directly affect melting and boiling points.

Key Terms

VSEPR (Valence Shell Electron Pair Repulsion)

A model that predicts the 3D arrangement of atoms around a central atom based on repulsion between bonding and nonbonding electron pairs in the valence shell. In simple terms, electron pairs push each other as far apart as possible, and that determines the molecule's shape.

Electron geometry (electronic geometry)

The spatial arrangement of all electron pairs (both bonding and nonbonding) around a central atom. Think of it as the shape you get when you count every electron group, including lone pairs.

Molecular geometry (molecular shape)

The spatial arrangement of only the bonded atoms, ignoring lone pairs. In simple terms, this is the shape you would see if you could photograph the molecule, because lone pairs are invisible.

Hybridisation

The mixing of atomic orbitals to form new, equivalent hybrid orbitals that accommodate bonding and lone pairs. sp = 2 groups, sp² = 3 groups, sp³ = 4 groups, sp³d = 5 groups, sp³d² = 6 groups. Think of it as the atom reshaping its orbitals to match however many electron groups surround it.

Sigma bond (σ bond)

A bond formed by head-to-head overlap of orbitals, with electron density concentrated along the axis between the two nuclei. In simple terms, the first bond between any two atoms is always a sigma bond, and it allows free rotation.

Pi bond (π bond)

A bond formed by side-by-side (parallel) overlap of p orbitals, with electron density above and below the bond axis. Think of it as the second or third bond in a double or triple bond. Pi bonds lock the atoms in place, preventing rotation.

Bonding orbital

A molecular orbital formed when atomic orbitals of the same sign (phase) combine, resulting in increased electron density between the nuclei and a lower-energy, more stable arrangement. In simple terms, it is the orbital that holds the atoms together.

Antibonding orbital

A molecular orbital formed when atomic orbitals of opposite sign combine, resulting in a node (region of zero electron density) between the nuclei and a higher-energy, less stable arrangement. Think of it as the orbital that would push atoms apart if electrons occupied it.

London dispersion forces (LDF)

Weak, short-lived intermolecular attractions caused by temporary, instantaneous dipoles in electron clouds. Present in all molecules. A type of Van der Waals force. In simple terms, the electrons in any atom briefly cluster on one side, creating a momentary imbalance that attracts neighbouring molecules.

Dipole-dipole interactions

Intermolecular attractions between the permanent partial charges of polar molecules. Think of it as polar molecules lining up so that each molecule's δ+ end faces another's δ− end.

Hydrogen bond

A particularly strong type of dipole-dipole interaction that occurs when hydrogen is bonded to a highly electronegative atom (F, O, or N) and is attracted to a lone pair on another F, O, or N atom. In simple terms, it is not a true bond (no electron sharing or transfer), but it is the strongest intermolecular force, roughly 10% as strong as a covalent bond.

Core Content

VSEPR Theory

• The 3D shape of a molecule is determined by repulsion between all electron pairs (bonding and nonbonding) around the central atom

• Lone pairs repel more strongly than bonding pairs, so they compress bond angles slightly

• Electron geometry counts all electron groups; molecular geometry counts only bonded atoms

Electron Geometry vs Molecular Geometry: Key Examples

• H₂O: four electron groups (two bonding pairs, two lone pairs). Electron geometry = tetrahedral. Molecular geometry = bent (angular). The two lone pairs push the bonding pairs closer together, reducing the bond angle below 109.5°.

• CH₄: four electron groups (all bonding). Electron geometry = tetrahedral. Molecular geometry = tetrahedral. No lone pairs, so bond angles are the ideal 109.5°.

• NH₃: four electron groups (three bonding, one lone pair). Electron geometry = tetrahedral. Molecular geometry = trigonal pyramidal. The lone pair compresses the bond angles slightly below 109.5°.

Polarity of Molecules

• A molecule with nonpolar bonds is always nonpolar

• A molecule with polar bonds can be either polar or nonpolar, depending on geometry

  • If the bond dipoles cancel due to symmetry, the molecule is nonpolar even though individual bonds are polar

  • Example: CH₂Cl₂ is nonpolar because the bond dipoles cancel in its tetrahedral geometry

  • Example: H₂O is polar because its bent shape means the dipoles do not cancel, producing a net dipole moment

Sigma and Pi Bonds

• Sigma bonds: formed by head-to-head orbital overlap. Electron density sits directly between the nuclei. Allow free rotation around the bond axis.

• Pi bonds: formed by parallel (side-by-side) overlap of unhybridised p orbitals. Electron density sits above and below the bond axis. No free rotation, because twisting would break the overlap.

• Bond composition:

  • Single bond = 1 σ

  • Double bond = 1 σ + 1 π

  • Triple bond = 1 σ + 2 π

Molecular Orbital Theory (Basics)

• When atoms bond, their atomic orbitals combine to form molecular orbitals that span both atoms

• If the orbital phases (signs) match, a bonding orbital forms (lower energy, more stable)

• If the phases are opposite, an antibonding orbital forms (higher energy, less stable, has a node between the nuclei)

Intermolecular Forces

Intermolecular forces are attractions between molecules (not within them). They directly determine melting points, boiling points, and solubility.

• London dispersion forces (weakest):

  • Present in all molecules, including nonpolar ones

  • Caused by temporary, instantaneous dipoles from random electron movement

  • Short-lived and do not extend over long distances

  • Larger, more polarisable molecules have stronger dispersion forces

• Dipole-dipole interactions (medium):

  • Present in polar molecules in solid and liquid phases; negligible in the gas phase

  • Polar molecules have higher boiling and melting points than nonpolar molecules of comparable molecular weight

  • London and dipole-dipole forces differ not in kind but in duration: both are electrostatic forces between partial charges, but dipole-dipole involves permanent rather than temporary dipoles

• Hydrogen bonds (strongest intermolecular force):

  • Occur when H is bonded to F, O, or N (remember the mnemonic: FON or FONS)

  • Can be intermolecular (between molecules) or intramolecular (within a molecule)

  • Not true bonds: no electron sharing or transfer occurs

  • Roughly 10% the strength of a typical covalent bond

  • Responsible for water's unusually high boiling point, surface tension, and role as a biological solvent

Hybridisation and Geometry Reference Table

Real-World Applications

• Water's bent geometry and resulting polarity make it the "universal solvent" for biological systems. Without that bent shape, water would be nonpolar and could not dissolve salts, sugars, or proteins, making life as we know it impossible.

• Hydrogen bonding between water molecules explains ice floating (ice is less dense than liquid water), which insulates aquatic ecosystems in winter and is fundamental to climate regulation.

Common Misconceptions

• Students often confuse electron geometry with molecular geometry. Electron geometry includes lone pairs; molecular geometry does not. H₂O has tetrahedral electron geometry but bent molecular geometry.

• A common error is assuming that any molecule with polar bonds is itself polar. Symmetry matters. If bond dipoles cancel (as in CO₂ or CH₂Cl₂), the molecule is nonpolar despite having polar bonds.

• Students sometimes think hydrogen bonds are intramolecular covalent bonds. They are not. "Hydrogen bond" is a misleading name for what is really a strong intermolecular attraction, with no electron sharing.

• Students frequently forget that London dispersion forces exist in all molecules, including polar ones. Polar molecules experience both London forces and dipole-dipole forces.

Why It Matters / Exam Flags

• ⚠️ VSEPR geometry predictions are heavily tested. Given a Lewis structure, you should be able to state the electron geometry, molecular geometry, hybridisation, and bond angles without hesitation.

• ⚠️ The distinction between electron geometry and molecular geometry is a classic exam question. If a molecule has lone pairs, those two geometries will differ.

• ⚠️ Sigma vs pi bond counting is a frequent short-answer item. Remember: single = 1σ, double = 1σ + 1π, triple = 1σ + 2π.

• ⚠️ Intermolecular forces ranking (London < dipole-dipole < hydrogen bonding) and the requirement for H bonded to F, O, or N for hydrogen bonds are consistently tested.

• ⚠️ Predicting whether a molecule is polar or nonpolar from its geometry is a common multi-step question combining Lewis structures, VSEPR, and dipole analysis.

Quick Self-Test

1. Fill in the blank: a molecule with four electron groups and two lone pairs has a ______ molecular geometry. (bent)

1. True or false: a double bond consists of two sigma bonds. (False, it is one sigma and one pi)

1. Fill in the blank: hydrogen bonding requires H to be bonded to ______, ______, or ______. (F, O, or N)

1. True or false: London dispersion forces only exist in nonpolar molecules. (False, they exist in all molecules)

1. Fill in the blank: sp³ hybridisation corresponds to ______ electron groups around the central atom. (four)

Practice Q&A

Q: What are the electron geometry and molecular geometry of water (H₂O)?

A: Electron geometry is tetrahedral (four electron groups: two bonding pairs plus two lone pairs). Molecular geometry is bent or angular, because the lone pairs are not visible in the shape but still repel the bonding pairs.

Q: Explain why CO₂ is nonpolar despite having two polar C=O bonds.

A: CO₂ is linear (two electron groups, no lone pairs on carbon). The two C=O bond dipoles point in exactly opposite directions and cancel each other, resulting in a net dipole moment of zero.

Q: How many sigma and pi bonds are in a molecule of ethylene (C₂H₄), which has a C=C double bond and four C–H single bonds?

A: The C=C double bond contributes 1σ + 1π. Each of the four C–H single bonds contributes 1σ. Total: 5 sigma bonds and 1 pi bond.

Q: Rank the following in order of increasing boiling point: CH₄, H₂O, HCl. Explain your reasoning.

A: CH₄ < HCl < H₂O. CH₄ is nonpolar, so it has only London dispersion forces (lowest boiling point). HCl is polar, adding dipole-dipole interactions. H₂O has hydrogen bonding (the strongest intermolecular force here), giving it the highest boiling point.

Q: What is the hybridisation of the central atom in SF₆, and what is its molecular geometry?

A: sp³d² hybridisation (six electron groups, all bonding). Molecular geometry is octahedral, with bond angles of 90°.

Connections to Other Topics

VSEPR geometry and polarity connect directly to intermolecular forces, which in turn explain bulk physical properties (boiling points, solubility, viscosity) covered in later chapters on states of matter. Molecular orbital theory is revisited in greater depth in organic chemistry, where sigma and pi bond character drives reaction mechanisms (additions, eliminations, substitutions).

Hydrogen bonding is central to biochemistry: it holds DNA's double helix together, stabilises protein secondary structures (alpha helices and beta sheets), and gives water the properties that sustain life.

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