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Intermolecular Forces

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Edexcel IAL Chemistry · Topic 2.2

Intermolecular Forces

The Big Idea: Molecules stick to their neighbours using weak "sticky" forces (much weaker than the covalent bonds holding atoms together inside them) — and the type and strength of these sticky forces explains almost everything about a substance's melting point, boiling point, and whether it dissolves in water.

Quick Summary

  • Intramolecular forces are inside a molecule (covalent bonds) — strong. Intermolecular forces are between molecules — weak.
  • There are three types of intermolecular force: induced dipole–dipole (London/van der Waals), permanent dipole–dipole, and hydrogen bonding (a special, extra-strong type of permanent dipole–dipole).
  • Strength order: Covalent bond > Hydrogen bond > Permanent dipole–dipole > Induced dipole–dipole.
  • Hydrogen bonding needs H directly bonded to N, O, or F (small, highly electronegative atoms with lone pairs).
  • Water's hydrogen bonds explain its unusually high boiling point, high surface tension, and why ice floats (is less dense than liquid water).
  • More electrons = stronger London forces = higher boiling point. Branching reduces surface contact = weaker London forces = lower boiling point.
  • "Like dissolves like": polar substances dissolve in polar solvents (water), non-polar substances dissolve in non-polar solvents.

1. Intramolecular vs Intermolecular Forces

Think of a box of magnets glued together in pairs. Each glued pair is stuck together really tightly — that's like a covalent bond, an intramolecular ("within the molecule") force. But when you tip the box, the pairs slide around and cling weakly to each other — that clinging is the intermolecular ("between molecules") force.

Intramolecular forces — forces within a molecule, holding its own atoms together. These are covalent bonds (single, double, triple, or co-ordinate/dative). They're formed when atoms share their outer electrons.
Intermolecular forces — weaker forces between neighbouring molecules. These are what you have to overcome to melt or boil a substance — you are NOT breaking covalent bonds when water boils, you're only breaking the hydrogen bonds between H₂O molecules.
INTERMOLECULAR force (weak, dashed) H-bond between two water molecules H H- - - - - -H H \ / \ / O O / / H H \___________/ INTRAMOLECULAR force (strong, solid line) the O-H covalent bond inside ONE molecule
Key Fact to Remember Intramolecular forces (covalent bonds) are roughly 10× stronger than intermolecular forces (like hydrogen bonds). This is WHY boiling water doesn't break it into H and O atoms — it just separates the molecules from each other.

Q1. When ice melts into liquid water, are covalent O–H bonds broken? Explain.

2. The Three Types of Intermolecular Force

All intermolecular forces are fundamentally about positive and negative charges attracting each other. The three types just differ in how permanent and how strong that charge separation is.

STRONGEST │ │ Hydrogen Bonding │ Permanent dipole – permanent dipole forces │ ▼ WEAKEST Induced dipole – induced dipole (London/van der Waals)
Analogy Time Imagine three ways two people could hold hands: (1) London forces = two strangers accidentally brushing hands in a crowd — brief, random, weak. (2) Permanent dipole–dipole = two people who always face each other and reach out — reliable and steady. (3) Hydrogen bonding = a firm handshake that only happens between very specific "VIP" people (N, O, F) — strongest of the three, but still nowhere near as strong as being surgically attached (covalent bond)!

2a. Induced Dipole–Dipole Forces (London Dispersion / van der Waals)

These exist between every single atom or molecule, polar or not — they're the one force that's always present, no exceptions. Here's the mechanism, step by step:

STEP 1: Electrons in a non-polar molecule are always moving. By random chance, more electrons end up on one side. ( ) --> (δ+ | δ-) an INSTANTANEOUS (temporary) dipole appears STEP 2: This temporary dipole repels/attracts the electron cloud of a NEIGHBOURING molecule, inducing a dipole in it too. (δ+ | δ-) ( ) --> (δ+ | δ-) (δ+ | δ-) temporary neutral the neighbour now has an dipole molecule INDUCED dipole STEP 3: The δ+ end of one molecule is attracted to the δ- end of its neighbour. This attraction = the intermolecular force. (δ+ | δ-) ‑ ‑ ‑ ‑ ‑ (δ+ | δ-) ^ this weak, constantly-flickering attraction is the London force

Because the electron cloud keeps moving, these dipoles are temporary — they appear and disappear constantly, which is why this force is so weak on its own.

Rule to Remember More electrons in a molecule → bigger, more frequent temporary dipoles → stronger London forces → higher melting/boiling point.

Q2. Why does xenon (Xe) have a higher boiling point than helium (He), even though neither is polar?

2b. Permanent Dipole–Dipole Forces

Some molecules are polar — because of differences in electronegativity between their atoms, they have a permanent δ+ end and a permanent δ− end (not a random, flickering one like London forces). Two polar molecules line up so the δ+ of one is attracted to the δ− of its neighbour.

POLAR MOLECULE POLAR MOLECULE δ+ δ- δ+ δ- (=======) - - - - - - (=======) permanent dipole – permanent dipole force (dashed line = intermolecular)
Worked Comparison — Butane vs Propanone Butane and propanone have the same number of electrons, so you might expect similar boiling points. But butane is non-polar (only London forces, weak), while propanone (CH₃COCH₃) is polar due to its C=O bond (has permanent dipole–dipole forces too). Result: butane boils at 0°C, but propanone boils at 56°C — proving permanent dipole–dipole forces are stronger than induced dipole forces alone, for molecules of similar size.

Q3. Two molecules have an identical number of electrons. One is polar, one is non-polar. Which has the higher boiling point, and why?

2c. Hydrogen Bonding — the "VIP" Intermolecular Force

Hydrogen bonding is a special, extra-strong type of permanent dipole–dipole force. It's the strongest intermolecular force, but it can only happen under very specific conditions.

Conditions for Hydrogen Bonding (both needed!) 1. A hydrogen atom covalently bonded to a very electronegative atom: N, O, or F.
2. That N/O/F atom must have a lone pair of electrons for the H to bond to on a neighbouring molecule.

Why only N, O, F? These three elements are small and highly electronegative, which means when they're bonded to hydrogen, they pull the shared electron pair strongly towards themselves. This leaves the H atom with a very large, concentrated δ+ charge (since hydrogen has no other electrons to shield its nucleus) — strong enough to be attracted right into a lone pair on a neighbouring N, O, or F atom.

δ+ δ- H₂O: H---O: ⋯⋯⋯⋯⋯ H---O ‥ (lone pair) δ+ ‥ (lone pair) H₂O Water (O has 2 lone pairs) → can form a MAXIMUM of 2 hydrogen bonds per molecule Ammonia NH₃ (N has only 1 lone pair) → can form a MAXIMUM of 1 hydrogen bond per molecule
Molecules that hydrogen bond Alcohols (O–H), water (O–H), ammonia (N–H), amines (N–H), carboxylic acids (O–H), hydrogen fluoride (H–F), and proteins (N–H). Notice the pattern — always look for H directly attached to N, O, or F.
Common Mistake Students often think ANY molecule with oxygen or nitrogen in it can hydrogen bond. NOT TRUE — the hydrogen must be directly, covalently bonded to the N, O, or F. For example, haloalkanes contain a polar C–X bond but CANNOT hydrogen bond, because there's no H–F, H–O or H–N bond present — the halogen isn't bonded to a hydrogen.

Q4. Explain why HF can hydrogen bond, but HCl cannot, even though both are hydrogen halides.

3. Water's Anomalous Properties

Water behaves "weirdly" compared to what you'd predict just from its tiny size — and hydrogen bonding is the reason for every single one of these anomalies.

High Melting & Boiling Points

If you plot enthalpy of vaporisation against number of electrons for the Group 16 hydrides (H₂O, H₂S, H₂Se, H₂Te), you'd expect H₂O — the smallest, with the fewest electrons — to have the lowest value, roughly 17 kJ/mol based on the trend from the other three. Instead, its actual enthalpy of vaporisation is almost 3× higher (~41 kJ/mol). That "anomalous jump" is the fingerprint of hydrogen bonding — an extra force present in H₂O that isn't present in H₂S, H₂Se or H₂Te (since S, Se and Te aren't electronegative enough to hydrogen bond).

Enthalpy of Vaporisation (kJ/mol) 50 | 40 | H2O ← anomalously high! | *. 30 | . | . 20 | *---*-------* | H2S H2Se H2Te 10 | +------------------------ 0 40 80 120 (electrons)

High Surface Tension

Molecules at the very surface of water are hydrogen-bonded to the molecules below and beside them, but not above (since there's air there). This uneven pulling drags surface molecules slightly inward/downward, compressing the surface into a kind of "skin." This is strong enough that some insects can walk on water!

Ice is Less Dense Than Liquid Water

This is the one that surprises people most — normally, solids are denser than their liquids because particles pack more tightly when frozen. Water does the opposite:

Why Ice Floats In ice, each water molecule is locked into a rigid, open lattice by hydrogen bonds — held at fixed distances and angles (the H–O–H bond angle even widens slightly, from 104.5° in liquid water to closer to the "ideal" 109.5° tetrahedral angle in ice). This hydrogen-bonded cage holds the molecules further apart than they'd be in the liquid, where they can pack more randomly and closely. Result: ice is about 9% less dense than liquid water — which is why ice cubes and icebergs float.

Q5. Predict and explain how the enthalpy of vaporisation of H₂O compares to H₂S if hydrogen bonding didn't exist — then explain why the real value is different.

4. Physical Properties & Intermolecular Forces

Branching Reduces Boiling Point

The larger the surface area of a molecule, the more contact points it has with neighbouring molecules, and the stronger the total London dispersion forces. Branching a molecule makes it more compact and spherical, reducing its surface area and contact with neighbours.

PENTANE (straight chain) 2,2-DIMETHYLPROPANE (branched) CH3-CH2-CH2-CH2-CH3 CH3 many contact points along | the chain = MORE surface area CH3--C--CH3 compact, spherical | shape = LESS surface Boiling point: 36°C CH3 area for contact Boiling point: 10°C
Rule to Remember Same number of electrons, but MORE branching → LESS surface contact → weaker London forces → LOWER boiling point.

Alcohols vs Alkanes

Alkanes only have induced dipole–dipole (London) forces. Alcohols have an O–H group, so they can ALSO hydrogen bond. This extra force means it takes much more energy to separate alcohol molecules.

Worked Example Propane (C₃H₈) boils at −42°C. Propanol (C₃H₇OH), similar size, boils at 97°C — nearly 140°C higher! The only structural difference is the O–H group enabling hydrogen bonding in propanol.
Common Mistake Don't say "alcohols have stronger van der Waals forces than alkanes" — that's not really the reason. Both still have similar London forces for similar-sized molecules. The real reason for the huge boiling point gap is the ADDITIONAL hydrogen bonding present only in the alcohol.

Q6. Hexanol (C₆H₁₃OH) is far less soluble in water than ethanol (C₂H₅OH), even though both contain an O–H group. Explain why.

5. Solvent Choice — "Like Dissolves Like"

Whether something dissolves depends on whether the solute's intermolecular forces can be effectively replaced by new forces with the solvent.

Non-polar solutes dissolve in non-polar solvents (e.g. hydrocarbons) — they form dispersion (London) forces with the solvent.
Polar covalent solutes dissolve in polar solvents (e.g. water) via dipole–dipole interactions or hydrogen bonds.
Ionic compounds dissolve in polar solvents — the polar solvent molecules surround and separate the ions, breaking down the ionic lattice.
ETHANOL dissolving in WATER — hydrogen bonds form between them: δ- δ+ H---O H H \ / \ / H---O--C---C--H : | (lone pair)H ethanol (CH3CH2OH) H bond water molecule

Why Ionic Compounds Dissolve in Water

Two-Step Mechanism 1. Polar water molecules surround and break down the ionic lattice.
2. The δ+ (hydrogen) ends of water surround negative anions; the δ− (oxygen) ends surround positive cations.

General rule: the greater the ionic charge, the LESS soluble the compound (stronger lattice = harder to break apart). E.g. 356.9 g of NaCl dissolves per dm³ of water, but only 74.4 g of CaCl₂ dissolves per dm³ — calcium's 2+ charge creates much stronger electrostatic attraction in the lattice, so it's harder for water to pull apart. (This is a general trend, not an absolute rule — there are exceptions.)

Haloalkanes — a Sneaky Exception

Watch Out! Haloalkanes contain a polar C–X bond (X = halogen), so you might expect them to dissolve well in water. But they only partially dissolve — because there's no H–F, H–O, or H–N bond present, so they CANNOT hydrogen bond with water. Polar bond ≠ automatic hydrogen bonding ability!

Q7. Predict whether iodine (I₂) would dissolve better in water or in hexane. Justify your answer.

What to Memorise

Term / RuleMeaning
Intramolecular forceForce WITHIN a molecule — covalent bonds (single/double/triple/co-ordinate). Strong.
Intermolecular forceForce BETWEEN molecules. Weak — roughly 1/10th the strength of a covalent bond.
Induced dipole–dipole (London/van der Waals)Weakest force; exists between ALL atoms/molecules; caused by momentary, random shifts in electron clouds.
Permanent dipole–dipoleAttraction between the fixed δ+ and δ− ends of polar molecules. Stronger than induced dipole forces for similar-sized molecules.
Hydrogen bondingStrongest intermolecular force. Needs H bonded directly to N, O, or F, attracted to a lone pair on N/O/F of a neighbouring molecule.
Strength orderCovalent bond > Hydrogen bond > Permanent dipole–dipole > Induced dipole–dipole (London)
Water: 2 H-bonds maxOxygen has 2 lone pairs → each H₂O molecule can form up to 2 hydrogen bonds.
Ammonia: 1 H-bond maxNitrogen has 1 lone pair → each NH₃ molecule can form up to 1 hydrogen bond.
More electrons→ stronger London forces → higher melting/boiling point.
More branching→ less surface area/contact → weaker London forces → lower melting/boiling point.
"Like dissolves like"Polar dissolves in polar; non-polar dissolves in non-polar; ionic compounds dissolve in polar solvents.
Ice is less dense than waterHydrogen bonds lock water molecules into an open, spaced-out lattice in ice — about 9% less dense than liquid.

Concepts Checklist

Exam Tips & Common Mistakes

Drawing Hydrogen Bonds Always use a dashed, straight line for intermolecular forces (never solid — that implies a covalent bond). The hydrogen bond must start at the lone pair (shown as dots) on the N/O/F atom and go right up to the δ+ hydrogen. Examiners specifically check that your dashed line starts and ends in the right place.
Trap: "Van der Waals" as a catch-all term Some students use "van der Waals forces" loosely to mean any intermolecular force. Be precise: van der Waals/London/induced dipole–dipole forces refer specifically to the WEAKEST type. If a question asks about a polar molecule's dominant force, don't just say "van der Waals" — name the correct, stronger force present (permanent dipole–dipole or hydrogen bonding).
Trap: Assuming polar bond = hydrogen bonding A polar bond alone is NOT enough for hydrogen bonding. You specifically need H bonded to N, O, or F. Haloalkanes (C–X polar bond) and things like CO₂ (polar bonds, but symmetric/no H bonded to O) do not hydrogen bond.
Comparing boiling points — always check TWO things When asked to explain a boiling point difference, examiners want you to check: (1) what TYPES of intermolecular force are present in each substance, and (2) if the same type is present in both, compare the number of electrons or degree of branching. Missing either step loses marks — don't just say "more van der Waals forces," explain WHY (more electrons, less branching, etc.)
Full-mark answer structure for "explain this trend" questions 1. Name the intermolecular force(s) involved.
2. Explain the underlying cause (e.g. more electrons → bigger temporary dipoles; O–H bond → hydrogen bonding).
3. State the consequence (more/less energy needed to separate molecules).
4. Link explicitly back to the property asked about (boiling point, solubility, density, etc.)
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Also in the full note
  • 4. Physical Properties & Intermolecular Forces
  • Exam Tips & Common Mistakes
  • High Melting & Boiling Points
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