Intermolecular Forces

Intermolecular forces (IM forces) are the comparatively weak attractions that act between discrete molecules. There are three to master at A-Level: London dispersion forces (sometimes called van der Waals forces); permanent dipole-dipole forces; and hydrogen bonding. These three differ in origin and in strength, and almost every physical property of a molecular substance — boiling point, melting point, viscosity, surface tension, solubility, enthalpy of vaporisation — is governed by which IM forces dominate and how many of them must be overcome. This lesson develops the strength hierarchy, applies it to the boiling-point trends of the Group 4–7 hydrides, and uses it to explain why water is the most anomalous small molecule in the periodic table — and why that anomaly is the chemical foundation of life on Earth.

Spec mapping (AQA 7405): This lesson anchors §3.1.3 (bonding — intermolecular forces). It cross-refers to lesson 3 of this course (electronegativity and bond polarity, §3.1.3, which drives permanent dipoles), to §3.1.4 (energetics — IM forces appear quantitatively in enthalpies of vaporisation and solution), to §3.3 (organic chemistry — boiling-point trends in alkanes, alcohols and carboxylic acids), and to §3.1.5 (kinetics — only collisions with sufficient kinetic energy can break IM forces in melting and boiling). Refer to the official AQA specification document for the exact wording of each section.

Assessment objectives: AO1 examines definitions and the three conditions for hydrogen bonding (H covalently bonded to N/O/F; lone pair on a neighbouring N/O/F; near-linear N–H···N or O–H···O alignment). AO2 expects students to predict the dominant IM force from a molecular structure and to rationalise boiling-point and density trends. AO3 demands explanation of water's anomalous physical properties (high mp/bp, density maximum at 4°C, high surface tension, high specific heat capacity) using hydrogen-bond network arguments — and connection of those properties to biological and environmental contexts.

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1. London Dispersion Forces (van der Waals)

London dispersion forces (also called dispersion forces, induced dipole-dipole forces, or simply van der Waals forces in much UK practice) are the universal intermolecular force. They act between every pair of molecules — polar or non-polar, large or small, gas, liquid or solid — wherever electrons exist.

Origin: instantaneous induced dipoles

The electron cloud of any molecule is not static. At any instant, by random fluctuation, the electrons can be slightly displaced to one side of the molecule, producing a transient instantaneous dipole with a partial negative charge on the electron-rich side and a partial positive charge on the electron-poor side. This instantaneous dipole electrostatically perturbs the electrons of a neighbouring molecule, inducing a complementary dipole oriented to give a net attraction:

   instantaneous            induced
        δ−   δ+              δ−   δ+
      [ A ]              [ B ]
              ←—— attraction ——→

The instantaneous-induced pair is short-lived (femtoseconds), but the fluctuation–induction process is constantly re-occurring across all pairs of nearby molecules; the time-averaged force is attractive and always present.

What governs the strength of dispersion forces?

Two factors dominate at A-Level:

1. Polarisability of the electron cloud. Polarisability scales with the number of electrons: a larger, more diffuse electron cloud distorts more easily. For molecules without a permanent dipole, dispersion-force strength tracks closely with relative molecular mass Mᵣ (because Mᵣ ≈ electron count in most common cases).
2. Molecular surface area / shape. Two molecules approach along a contact patch. A long, linear molecule presents more surface for contact than a compact, spherical one of the same Mᵣ, so its dispersion forces are larger.

Worked example: noble-gas boiling points

Noble gas	Mᵣ	Electrons	Bp (°C)
He	4	2	−269
Ne	20	10	−246
Ar	40	18	−186
Kr	84	36	−153
Xe	131	54	−108

Noble gases are monatomic and non-polar. Dispersion is the only IM force available. Boiling points climb monotonically down the group: more electrons → more polarisable atom → stronger dispersion → more energy needed to break the IM lattice → higher bp. This is the cleanest experimental fingerprint of dispersion forces in the entire periodic table.

Worked example: linear vs branched pentane

Two C₅H₁₂ isomers, identical Mᵣ (72) and identical electron count:

Isomer	Shape	Bp (°C)
Pentane (n-pentane)	linear chain	36
2,2-dimethylpropane (neopentane)	near-spherical	9.5

The molecular formula, mass and electron count are identical, so the only variable is shape. The linear chain offers a long contact strip where dispersion forces can act along the whole length; the spherical neopentane minimises contact area to a point. The 26.5°C bp difference is therefore entirely a surface-area effect on dispersion forces.

Boiling-point ladder for the alkanes

Alkane	Formula	Mᵣ	Bp (°C)
Methane	CH₄	16	−162
Ethane	C₂H₆	30	−89
Propane	C₃H₈	44	−42
Butane	C₄H₁₀	58	0
Pentane	C₅H₁₂	72	36
Hexane	C₆H₁₄	86	69
Decane	C₁₀H₂₂	142	174

Each added –CH₂– unit adds ~14 to Mᵣ and ~25–30°C to bp. Alkanes are non-polar (C and H have similar electronegativities), so dispersion is again the only IM force; the trend is essentially a calibration curve for dispersion strength versus Mᵣ.

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2. Permanent Dipole-Dipole Forces

Permanent dipole-dipole forces act between molecules that carry a permanent dipole moment — i.e. polar molecules in which the centres of positive and negative charge do not coincide because of electronegativity differences (see lesson 3).

Origin

In a polar molecule the bonded electrons sit closer to the more electronegative atom, giving that atom a partial negative charge (δ−) and the less electronegative atom a partial positive charge (δ+). When the overall geometry does not cancel these bond dipoles, the molecule has a net dipole moment. Nearby polar molecules orient so that the δ+ region of one faces the δ− region of another:

    δ+  δ−       δ+  δ−       δ+  δ−
    H—Cl    ···   H—Cl    ···   H—Cl

The electrostatic interaction is attractive on average across thermal motion in the liquid.

Strength

For molecules of similar Mᵣ, permanent dipole-dipole forces add on top of the dispersion forces, raising the total IM strength. Typical magnitudes are 5–20 kJ mol⁻¹ per pairwise interaction, compared with 1–10 kJ mol⁻¹ for dispersion in small molecules. Crucially, dispersion does not switch off when a molecule is polar — it is always present; permanent dipole-dipole is an addition, not a replacement.

Worked example: HCl, CH₃Cl, propanone

Molecule	Mᵣ	Dipole moment / D	Bp (°C)
HCl	36.5	1.05	−85
CH₃Cl (chloromethane)	50.5	1.87	−24
(CH₃)₂CO (propanone)	58	2.88	56

Worked example: propanone vs butane (same Mᵣ)

Propanone (CH₃COCH₃, Mᵣ = 58, bp = 56°C) and butane (C₄H₁₀, Mᵣ = 58, bp = 0°C) have identical Mᵣ and similar electron counts. Butane is non-polar; propanone has a strong C=O dipole. The 56°C bp gap is the unambiguous signature of permanent dipole-dipole forces on top of dispersion.

Exam Tip: In any polar molecule, dispersion is usually still the largest single component of the total IM force; permanent dipole-dipole is the differentiator that explains why polar molecules boil at higher temperatures than non-polar molecules of similar Mᵣ.

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3. Hydrogen Bonding

Hydrogen bonding is a uniquely strong subclass of permanent dipole-dipole interaction. It is responsible for water's anomalous properties, the double-helix architecture of DNA, the secondary structure of proteins, and the solubility of alcohols and sugars in water — to name only the obvious cases.

Three conditions (the AQA AO1 checklist)

A hydrogen bond requires all three of the following:

1. A hydrogen atom covalently bonded to N, O, or F (the H-bond donor).
2. A lone pair on a neighbouring N, O, or F atom (the H-bond acceptor).
3. A near-linear N–H···N, O–H···O, N–H···O, F–H···F (or other N/O/F combination) alignment. The lone pair must point along the H-X axis for maximum overlap.

    δ+        δ−       δ+        δ−
    —O—H ····· :O—H ····· :O—
              hydrogen bond

Why only N, O, F?

Hydrogen bonding is the strong-end limit of dipole-dipole interaction. Three features of N, O and F combine to produce it:

• High electronegativity (N 3.0, O 3.5, F 4.0) places a large δ+ on the hydrogen.
• Small atomic radius concentrates the lone pair in a small volume close to the nucleus, allowing close approach of the δ+ hydrogen.
• Lone pair geometry — N, O, F all have at least one available lone pair pointing outward from the atom.

Chlorine has electronegativity 3.0 (equal to N) but is much larger; its lone pairs are diffuse and far from the nucleus, so it does not form classical hydrogen bonds in A-Level treatment.

Strength

Hydrogen bonds are typically 10–40 kJ mol⁻¹ per bond — roughly 10× stronger than permanent dipole-dipole, and ~1/10 the strength of a typical covalent bond. The key hierarchy at A-Level:

IM force	Typical energy / kJ mol⁻¹
London dispersion (small molecules)	1–10
Permanent dipole-dipole	5–20
Hydrogen bonding	10–40
(Covalent bond, for comparison)	200–500
(Ionic lattice, for comparison)	600–4000

This is why boiling a covalent molecular substance breaks IM forces (low energy, low bp) but boiling sodium chloride breaks ionic bonds (very high energy, very high bp).

IM forces vs covalent and ionic bonds

IM forces are inter-molecular — between separate molecules. Covalent bonds are intra-molecular — within a molecule. Ionic bonds operate within an ionic lattice. When water boils at 100°C, the H–O covalent bonds (464 kJ mol⁻¹) are not broken; only the hydrogen bonds between water molecules (~23 kJ mol⁻¹) are broken. Conflating these is the single most common A-Level error in this topic.

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4. Anomalous Boiling Points of the Group 4–7 Hydrides

A famous data set — and the cleanest case study of hydrogen bonding at A-Level — is the bp of the simple hydrides as you descend each of Groups 4, 5, 6 and 7:

Period 2	Bp / °C	Period 3	Bp / °C	Period 4	Bp / °C	Period 5	Bp / °C
CH₄	−162	SiH₄	−112	GeH₄	−88	SnH₄	−52
NH₃	−33	PH₃	−87	AsH₃	−62	SbH₃	−17
H₂O	+100	H₂S	−60	H₂Se	−41	H₂Te	−2
HF	+20	HCl	−85	HBr	−67	HI	−35

Three observations:

• Group 4 (CH₄, SiH₄, GeH₄, SnH₄): smooth rise with Mᵣ — pure dispersion-force trend, no hydrogen bonding (C is not electronegative enough).
• Groups 5, 6, 7: the second-period member is dramatically anomalously high — NH₃, H₂O, HF all stand far above the dispersion-only extrapolation from their heavier congeners.
• The size of the anomaly is largest for H₂O (water bp is ~180°C higher than dispersion alone would predict), then HF, then NH₃ — tracking the strength and number of available H-bonds per molecule.

Why is water's anomaly the largest?

H₂O has two O–H bonds (two H-bond donors) and two oxygen lone pairs (two H-bond acceptors). Each water molecule can therefore participate in up to four hydrogen bonds simultaneously, forming a three-dimensional H-bond network. HF has one H but three lone pairs (donor-limited, max 2 H-bonds per molecule); NH₃ has three H but only one lone pair (acceptor-limited, max 2 H-bonds per molecule). Water is the only one balanced for full 3D networking — hence its outsized anomaly.

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5. Water — the Anomalous Liquid

Water's H-bond network produces a cluster of macroscopic anomalies, every one of which is biologically and environmentally consequential.

(a) Anomalously high boiling and melting point

H₂O bp = 100°C and mp = 0°C. Extrapolation from the heavier Group 6 hydrides predicts bp ≈ −80°C. Without hydrogen bonding water would be a gas at typical Earth-surface temperatures — and no liquid-water-dependent biology could exist.

(b) Density maximum at 4°C (ice less dense than water)

Most liquids contract continuously on cooling and the solid is denser than the liquid. Water reverses this:

• From 100°C down to 4°C, water contracts (density rises).
• At 4°C, density reaches a maximum of 1.000 g cm⁻³.
• From 4°C down to 0°C, water expands slightly (density falls).
• On freezing, water expands further: ice has density 0.917 g cm⁻³, about 9% less than liquid water.

The molecular explanation is that ice forms an open hexagonal lattice in which every water molecule is held at fixed tetrahedral geometry by four hydrogen bonds. The lattice contains large hexagonal voids — wasted volume. On melting, ~10–15% of these H-bonds break, the lattice partially collapses, and the molecules pack more efficiently (denser). Between 0 and 4°C, two competing effects act: thermal expansion (normal) and progressive lattice collapse (extra contraction); below 4°C, lattice collapse dominates; above 4°C, thermal expansion dominates. The crossover is the 4°C density maximum.

Environmental consequence: ice floats on lakes and oceans, insulating the liquid water beneath. Aquatic life survives winter only because of this anomaly.

(c) High surface tension

Water has a surface tension of 72 mN m⁻¹ at 25°C — among the highest of any liquid except mercury (a metallic case). Surface molecules can H-bond only sideways and downward (not upward, into air), giving rise to a strong net inward pull. Consequences: capillary action, the meniscus, raindrops being spherical, water striders walking on water.

(d) High specific heat capacity

Water's specific heat capacity is 4.18 J g⁻¹ K⁻¹, roughly twice that of most other liquids. Heating water requires energy not only to increase the kinetic energy of the molecules but also to partially break the H-bond network (the rotational and translational freedom of water molecules is restricted by H-bonds at lower temperatures and increases as T rises). Hence a large fraction of input heat goes into H-bond reorganisation rather than translational kinetic energy. Consequence: oceans buffer global climate dramatically — they absorb and release enormous quantities of heat with small temperature change.

(e) High enthalpy of vaporisation

ΔH_vap(H₂O) = +41 kJ mol⁻¹ — anomalously large for such a small molecule (compare CH₄ at 8 kJ mol⁻¹). This is the macroscopic measure of how much energy is required to break the H-bond network on going from liquid to gas. Consequence: sweat cooling in humans, and the latent-heat-driven mechanics of weather systems.

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6. Hydrogen Bonding in Biology

Hydrogen bonding is the structural glue of the living world. A-Level Chemistry signposts but does not exhaust the following examples — they are picked up properly in §3.3.13 (amino acids, proteins, DNA).

• DNA base pairing. Adenine pairs with thymine via 2 hydrogen bonds; guanine pairs with cytosine via 3 hydrogen bonds. These bonds are individually weak but collectively hold the double helix together. They are weak enough to be unzipped on demand during replication and transcription — a property absolutely essential to the function of the molecule.
• Protein secondary structure. The α-helix and β-sheet are stabilised by hydrogen bonds between C=O and N–H groups of the peptide backbone. Denaturation of a protein (by heat or pH) is principally the breaking of this H-bond pattern.
• Cellulose. Long parallel chains of poly-β-glucose are held in rigid sheets by interchain hydrogen bonds — the basis of plant cell-wall strength and a major reason cellulose is hard for animal enzymes to digest.
• Phospholipid bilayers. The polar head groups of phospholipids hydrogen-bond with surrounding water, anchoring the bilayer architecture of every cell membrane.
• Alcohol and sugar solubility. Ethanol and glucose dissolve in water because their O–H groups can act as both H-bond donors and acceptors; the H-bonding network of water is reorganised, not disrupted.

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7. The Organic Bp Ladder — Alkane, Alcohol, Carboxylic Acid

A useful synoptic case (§3.3 onwards) is the bp progression for organic molecules of similar Mᵣ:

Molecule	Formula	Mᵣ	IM forces	Bp (°C)
Propane	CH₃CH₂CH₃	44	dispersion only	−42
Ethanal	CH₃CHO	44	dispersion + dipole-dipole	20
Ethanol	CH₃CH₂OH	46	dispersion + dipole-dipole + H-bonding	78
Methanoic acid	HCOOH	46	dispersion + dipole-dipole + H-bonding (dimer)	101

The bp ladder is a direct read-out of the IM-strength hierarchy: dispersion-only at the bottom, H-bonding upper, and carboxylic-acid dimerisation (two H-bonds between the COOH groups of paired molecules) at the top.

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8. Summary Table — Predicting the Dominant IM Force