Intermolecular Forces

Spec Mapping — OCR H432 Module 2.2.2 — Bonding and structure, covering the three categories of intermolecular force (London dispersion forces, permanent dipole-dipole forces, hydrogen bonding), the origin and strength of each, identification of which forces operate in a given molecule, and the use of intermolecular forces to explain physical properties — particularly boiling-point trends across the hydrides of Groups 14, 15, 16, 17 and within homologous series of organic molecules (refer to the official OCR H432 specification document for exact wording).

Intermolecular forces are the attractive interactions between separate molecules. They are distinct from — and much weaker than — the covalent bonds within molecules. When a molecular substance melts or boils, you break intermolecular forces, not covalent bonds. The strength and identity of the intermolecular forces between molecules determine essentially every observable physical property: melting point, boiling point, solubility, viscosity, surface tension, vapour pressure. This lesson categorises the three intermolecular forces examined at OCR A-Level — London dispersion forces, permanent dipole-dipole forces, and hydrogen bonding — explains their origin, ranks them by strength, and applies them to explain the canonical OCR trends in hydride boiling points. The lesson sits at the centre of the bonding module because every prior lesson (electronegativity, dipole moment, molecular shape) feeds into the intermolecular force analysis, and every subsequent topic in chemistry uses intermolecular forces as the bridge from molecular structure to bulk behaviour.

Key Definitions:

• Intermolecular force (IMF) — an attractive force acting between separate molecules (as distinct from the intramolecular covalent bonds within a molecule).
• London dispersion force (also: induced dipole-dipole force, van der Waals force) — a weak attraction between an instantaneous dipole on one molecule and the induced dipole it creates in a neighbouring molecule.
• Permanent dipole-dipole force — an attraction between the δ⁺ end of one polar molecule and the δ⁻ end of a neighbouring polar molecule.
• Hydrogen bond — a particularly strong dipole-dipole interaction between a δ⁺ H bonded to N, O, or F and a lone pair on another N, O, or F.

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Why intermolecular forces matter

Covalent bonds (typically 150–1000 kJ mol⁻¹) hold atoms together within a molecule. Intermolecular forces (typically 1–40 kJ mol⁻¹) hold molecules together in liquid or solid phase. When you heat water from 25 °C to 100 °C and boil it, you are providing the ~44 kJ mol⁻¹ enthalpy of vaporisation to break the intermolecular forces between water molecules — but you are not breaking the O-H covalent bonds (which would require ~460 kJ mol⁻¹). The product of boiling is water molecules in the gas phase, not atomic O and H.

This 10-100× ratio between covalent bond strength and intermolecular force strength is the reason molecular substances have much lower melting and boiling points than ionic compounds (where you must break the entire ionic lattice to melt) or giant covalent structures (where you must break covalent bonds to melt).

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The three types of intermolecular force — overview

flowchart TD
    A[Intermolecular forces between molecules] --> B[London dispersion forces]
    A --> C[Permanent dipole-dipole forces]
    A --> D[Hydrogen bonding]
    B --> B1[Present in ALL substances]
    B --> B2[Strength: 1-10 kJ mol-1 typical]
    B --> B3[Grows with electron count and surface area]
    C --> C1[Polar molecules only]
    C --> C2[Strength: 5-25 kJ mol-1]
    C --> C3[Operates IN ADDITION to London forces]
    D --> D1[H bonded to N O F only]
    D --> D2[Strength: 10-40 kJ mol-1]
    D --> D3[Operates IN ADDITION to London + dipole-dipole]

Typical strength comparison

Force type	Typical strength (kJ mol⁻¹)	Required conditions
London dispersion	1 – 10 (much more for large molecules)	None — present in all matter
Permanent dipole-dipole	5 – 25	Polar molecule
Hydrogen bonding	10 – 40	H bonded to N, O, or F + N/O/F lone pair acceptor
Covalent bond (for comparison)	150 – 1000+	Two atoms sharing electrons
Ionic lattice (for comparison)	600 – 4000+	Ions in a lattice

Crucially, intermolecular forces add up cumulatively: a polar molecule with H bonded to O experiences London forces plus dipole-dipole plus hydrogen bonding — all three at once, not one to the exclusion of others.

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1. London dispersion forces

Named after Fritz London (1930), who derived a quantum-mechanical description. Also called induced dipole-dipole, instantaneous dipole-induced dipole, or van der Waals forces (though the latter strictly refers to a broader category including dipole-dipole).

Origin

At any instant the electrons in a molecule are not perfectly symmetrically distributed — random thermal fluctuations create a brief instantaneous dipole. This tiny transient dipole electrostatically distorts the electron cloud of a neighbouring molecule, inducing a corresponding instantaneous dipole. The two dipoles attract for the duration of the fluctuation.

Over time the instantaneous dipoles continually appear, vanish, and reorient — but the induced response always lags slightly and always corresponds to attraction (never repulsion), so the net effect averaged over time is an attractive force between all molecules. London dispersion forces are therefore always present between any two molecules.

Key features

• Present in all molecules (polar or non-polar; covalent or ionic).
• The only intermolecular force between non-polar molecules.
• Strength depends on:
  • Number of electrons (≈ molar mass): more electrons mean more polarisable electron clouds and stronger instantaneous dipoles.
  • Surface area of molecular contact: straight-chain molecules pack closer than branched isomers, increasing the available area for transient dipole interactions.

Trend example — straight-chain alkanes

Alkane	M_r	Boiling point (°C)
CH₄	16	−162
C₂H₆	30	−89
C₃H₈	44	−42
C₄H₁₀	58	−0.5
C₈H₁₈	114	126
C₁₆H₃₄	226	287

Each step up the homologous series adds ~14 e⁻ and ~30 °C to the boiling point — a clean demonstration that London force strength scales with electron count.

Trend example — Group 17 halogens

Halogen	Electrons / molecule	State at 25 °C	Boiling point (°C)
F₂	18	Gas	−188
Cl₂	34	Gas	−34
Br₂	70	Liquid	59
I₂	106	Solid	184

All non-polar; only London forces between molecules. Increasing electron count → stronger London forces → higher b.p. and progressively more condensed phase at room temperature.

Branching effect — pentane isomers

Isomer	Structure	Boiling point (°C)
n-pentane	straight chain	36
2-methylbutane	one branch	28
2,2-dimethylpropane (neopentane)	fully branched (almost spherical)	9.5

All three have the same molecular formula (C₅H₁₂) and same Mr (72), so London force differences come entirely from surface contact area. Branched (compact) isomers have less surface contact and weaker London forces.

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2. Permanent dipole-dipole forces

Origin

Polar molecules (with a permanent dipole moment — see Lesson 9) have δ⁺ and δ⁻ ends. The δ⁺ end of one molecule electrostatically attracts the δ⁻ end of a neighbouring molecule. Unlike London forces (transient), permanent dipole-dipole forces are continuous and oriented.

δ⁺ δ⁻ δ⁺ δ⁻ H Cl H Cl Dipole-dipole attraction (δ⁻ of one Cl atom attracts δ⁺ of next H atom)

Typical dipole-dipole strength: 5–25 kJ mol⁻¹ Operates IN ADDITION to ever-present London forces

Key features

• Present only in polar molecules (those with a net dipole moment per Lesson 9).
• Operates in addition to London forces (always cumulative).
• Strength: 5–25 kJ mol⁻¹, depending on dipole moment magnitude.

Comparison example — HCl vs F₂

Property	HCl	F₂
Molar mass	36.5	38
Electrons / molecule	18	18
Polar?	Yes (μ = 1.08 D)	No (homonuclear)
Forces present	London + dipole-dipole	London only
Boiling point	−85 °C	−188 °C

Similar Mr and electron count → similar London forces. The additional 5–10 kJ mol⁻¹ of dipole-dipole interactions in HCl raise its boiling point by ~100 °C. The HCl/F₂ comparison is the canonical OCR example for showing that dipole-dipole > London alone.

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3. Hydrogen bonding

Origin and conditions

Hydrogen bonding is a particularly strong dipole-dipole interaction that arises when H is covalently bonded to one of three small, highly electronegative atoms — N, O, or F — and is then attracted to a lone pair on another N, O, or F atom of a neighbouring molecule.

The H atom becomes unusually δ⁺ because:

1. The bonding pair is pulled strongly toward the highly electronegative N/O/F.
2. H has no inner electrons to shield its nucleus — once the bonding pair is pulled away, the bare proton-like nucleus is exposed.
3. H is very small (atomic radius ~25 pm), allowing it to approach a lone pair on a neighbouring atom very closely.

The combination of large δ⁺ + small radius + accessible lone pair on the acceptor produces an unusually strong electrostatic interaction — typically 20 kJ mol⁻¹, ten times a typical London force.

O H H ·· δ⁻ δ⁺

Hydrogen bond (H···O lone pair)

O ·· H H

O ·· H

Requirements: • H covalently bonded to N, O, or F • Acceptor lone pair on N, O, or F • Linear D-H···A geometry (~180°) • Strength: 10-40 kJ mol⁻¹ • Show H-bond as DASHED line

The N/O/F-only rule

Hydrogen bonding requires both conditions simultaneously:

1. A H atom covalently bonded to N, O, or F (a δ⁺ donor, often written D-H).
2. A lone pair on another N, O, or F atom (an acceptor, A).

H bonded to C does not form hydrogen bonds because Δχ(C-H) = 0.35 is too small to generate a sufficient δ⁺ on H. Similarly, H bonded to Cl, S, P, or any other element fails the test — H must be bonded to N, O, or F specifically. This sharp rule is one of the most-tested OCR mark scheme details.

Strength and geometry

A hydrogen bond is typically 20 kJ mol⁻¹ (range 10–40 kJ mol⁻¹), much weaker than an O-H covalent bond (460 kJ mol⁻¹) but far stronger than ordinary dipole-dipole forces. Hydrogen bonds prefer linear geometry (D-H···A angle ~180°) for maximum orbital overlap between the H 1s and the acceptor lone pair.

Examples of hydrogen-bonding substances

• H₂O, ice, alcohols — O-H bonded H + O lone pair acceptor.
• NH₃, amines — N-H bonded H + N lone pair acceptor.
• HF — F-H bonded H + F lone pair acceptor.
• Carboxylic acids (e.g. CH₃COOH) — form dimers via two simultaneous H-bonds, giving anomalously high b.p.
• Proteins — α-helix and β-sheet secondary structures held by N-H···O=C backbone H-bonds.
• DNA double helix — base pairs held by 2 H-bonds (A-T) or 3 H-bonds (G-C).

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Identifying intermolecular forces in a substance

Always check in this order:

1. London forces — always present.
2. Permanent dipole-dipole — add if the molecule is polar (has net dipole moment).
3. Hydrogen bonding — add if the molecule has H bonded to N, O, or F and another N/O/F atom present (with lone pair) for the acceptor.

Molecule	London	Dipole-dipole	H-bond	Total
CH₄	Yes	No (non-polar)	No (H on C)	London only
CO₂	Yes	No (linear, cancels)	No	London only
HCl	Yes	Yes	No (H on Cl)	London + dipole
CHCl₃	Yes	Yes	No (H on C)	London + dipole
NH₃	Yes	Yes	Yes	All three
H₂O	Yes	Yes	Yes	All three
CH₃OH	Yes	Yes	Yes	All three
CH₃COOH	Yes	Yes	Yes (dimer)	All three
CH₃F	Yes	Yes	No (H on C, not F)	London + dipole

The CH₃F case is notable: there is F in the molecule, but no H is bonded to F (the F is bonded to C, while the H atoms are bonded to C). Therefore no hydrogen bonding.

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