Teaching Intermolecular Forces: A Lesson Guide

Ask a class what happens to water molecules when a kettle boils, and a good number will tell you the molecules “split into hydrogen and oxygen”. That single answer shows why intermolecular forces are hard to teach. Students have spent months learning that bonds hold atoms together, and now they must accept a second, much weaker kind of attraction acting between whole molecules. If the difference between the two is not nailed down early, every later explanation of boiling points, solubility and viscosity rests on a mistake.

This guide sets out a three-lesson sequence for students of about 15 to 17, with timings, low-risk practical work, a data-analysis task built on real boiling points, and the misconceptions you are almost certain to meet. It assumes you teach the three standard categories: London (dispersion) forces, permanent dipole–dipole forces and hydrogen bonds.

Learning goals

By the end of the sequence students should be able to:

1. Distinguish intramolecular bonds (covalent bonds within a molecule) from intermolecular forces (attractions between molecules).
2. Explain how London forces arise from temporary dipoles and why they grow with the number of electrons and with surface contact.
3. Identify polar molecules and describe permanent dipole–dipole attractions.
4. State the conditions for hydrogen bonding and draw a correctly labelled hydrogen bond.
5. Use intermolecular forces to explain trends in boiling point, volatility and solubility.

Prior knowledge to check

• Covalent bonding and dot-and-cross diagrams for simple molecules.
• Molecular shapes (at least tetrahedral, pyramidal, bent and linear). Polarity depends on shape, so if VSEPR hasn’t been taught yet, teach it first (see teaching molecular shapes with VSEPR).
• Electronegativity and polar bonds.
• The particle model of solids, liquids and gases, and the idea that changing state means particles separating.

A five-question diagnostic at the start of lesson 1 is worth the time. Include the kettle question above. The answers tell you how much untangling lesson 1 has to do.

Lesson 1: bonds versus forces (60 minutes)

Starter: what breaks when water boils? (10 minutes)

Show a photograph of a boiling kettle and ask students to draw what the particles look like in the steam. Collect the drawings without comment. Then show two numbers side by side: roughly 460 kJ to break one mole of O–H bonds, compared with about 41 kJ to boil one mole of water. Ask: “If boiling broke O–H bonds, would a kettle manage it?” The contrast in scale does most of the work.

Main: intramolecular versus intermolecular (20 minutes)

Build two water molecules from a model kit and join them loosely with a short elastic band or a strip of sticky tack between the O of one and an H of the other. The rigid sticks are covalent bonds; the floppy band is the intermolecular attraction. Pull the molecules apart: the band gives way long before any stick does.

Get students to write a two-column definition, “bond within a molecule” and “force between molecules”, and give three examples of each change: boiling ethanol, melting ice and dissolving sugar in one column, burning methane and electrolysing water in the other.

Activity: the evaporation race (20 minutes)

Place single drops of water, ethanol and propanone on a paper towel (or on the back of the hand, using only water and ethanol) and time how long each takes to disappear. Propanone and ethanol are flammable: keep all flames out of the room, use dropper bottles containing only a few cm³, and ventilate. Students rank the liquids by volatility and suggest which has the strongest attractions between molecules. Don’t explain the ranking yet. Lesson 3 comes back to it with the full model.

Plenary (10 minutes)

Hand back the steam drawings. Students correct their own, adding a sentence: “When water boils, the ___ between molecules are overcome; the ___ bonds are not broken.”

Lesson 2: the three kinds of force (60 minutes)

London forces (20 minutes)

Use a crowd analogy: electrons in a molecule are constantly moving, and at any instant there may be more on one side than the other. That creates a temporary dipole, which induces a dipole in the next molecule, and the two attract. Stress three points:

• London forces act between all molecules, polar or not.
• They grow with the number of electrons (bigger electron clouds are more easily distorted).
• They grow with surface contact: long, unbranched molecules touch along more of their length.

Give students the alkane boiling points and ask them to plot boiling point against the number of carbon atoms:

Alkane	Boiling point (°C, approximate)
Methane, CH₄	−162
Ethane, C₂H₆	−89
Propane, C₃H₈	−42
Butane, C₄H₁₀	−1
Pentane, C₅H₁₂	36
Hexane, C₆H₁₄	69

Follow up with pentane (36 °C) against its branched isomer 2,2-dimethylpropane (about 10 °C): same formula, same number of electrons, different shape. This pair is the clearest evidence that contact area matters. More detail is in London dispersion forces.

Permanent dipole–dipole forces (15 minutes)

Revisit polarity with molecular models: a molecule is polar if it has polar bonds and a shape in which the bond dipoles don’t cancel. CO₂ has polar bonds but is linear, so it is non-polar; H₂O is bent and polar. Molecules with permanent dipoles line up so that δ+ ends sit near δ− ends, in addition to their London forces. See dipole–dipole forces.

Hydrogen bonding (20 minutes)

Set out the three conditions on the board:

1. A hydrogen atom covalently bonded to N, O or F.
2. A lone pair on an N, O or F atom in a neighbouring molecule.
3. The hydrogen bond forms between that H (strongly δ+) and the lone pair, in a roughly straight line O–H···O.

Have students draw hydrogen bonds between two water molecules, then between ammonia and water, using a dashed line and labelling δ+, δ− and the lone pair. Insist the dashed line goes to the lone pair, not to the hydrogen of the other molecule. Hydrogen bonding has more worked diagrams.

Plenary (5 minutes)

Card sort: ten molecules (CH₄, HCl, H₂O, NH₃, CO₂, CH₃OH, I₂, HF, CH₃Cl, C₂H₆). Students put each into every category that applies. Methanol should appear under all three.

Lesson 3: explaining the data (60 minutes)

Data task: the hydride boiling-point graph (25 minutes)

Give groups the boiling points of the hydrides of groups 14 to 17 and ask them to plot each group as a line against period:

Group 14	b.p. (°C)	Group 15	b.p. (°C)	Group 16	b.p. (°C)	Group 17	b.p. (°C)
CH₄	−162	NH₃	−33	H₂O	100	HF	20
SiH₄	−112	PH₃	−88	H₂S	−60	HCl	−85

(Extend with the period 4 and 5 hydrides if you have a data book.) Questions to ask as they work:

• Why does the group 14 line rise steadily? (More electrons, stronger London forces; no hydrogen bonding in any of them.)
• Which three points “break the pattern”? (NH₃, H₂O, HF.)
• Why is water’s boiling point higher than that of HF, even though F is more electronegative? (Each water molecule has two H atoms and two lone pairs, so it can form on average two hydrogen bonds per molecule in the liquid; HF has three lone pairs but only one H.)

Why water has a high boiling point is a good follow-up reading for homework.

Return to the evaporation race (10 minutes)

Students now explain the lesson 1 ranking. Water evaporated slowest because it forms the most extensive hydrogen bonding. Ethanol forms hydrogen bonds too, but has only one O–H per molecule. Propanone has no O–H, so it relies on dipole–dipole and London forces and is the most volatile.

Demonstration: surface tension (10 minutes)

Float a paperclip on the surface of water in a clean Petri dish, lowering it gently on a small piece of tissue paper that then sinks. Add one drop of washing-up liquid at the edge and the paperclip sinks. Ask students to explain in terms of forces between water molecules at the surface and how the detergent interferes with them. There are no hazards beyond spilled water.

Exit questions (15 minutes)

See the question bank below; choose four, mixing recall and explanation.

Misconceptions to expect

Misconception	What to do
“Boiling breaks covalent bonds.”	The energy comparison (460 vs 41 kJ); kit models with sticks and elastic bands; always ask “which forces are overcome?”
“A hydrogen bond is a bond to a hydrogen atom.”	Call it “a hydrogen bond between molecules”; compare it with the O–H covalent bond in the same diagram
“Hydrogen bonds form with any H.”	CH₄ contains hydrogen but cannot form hydrogen bonds; test with C–H molecules in the card sort
“Polar bonds mean a polar molecule.”	CO₂ and CCl₄: symmetric shapes cancel the dipoles
“Non-polar molecules have no intermolecular forces.”	Iodine is a solid at room temperature with only London forces
“London forces are always the weakest.”	Compare hexane (69 °C, London only) with the polar hydrogen chloride (−85 °C)
“Hydrogen bonds are stronger than covalent bonds.”	They are roughly a tenth to a twentieth as strong, or less. Use the energy numbers again

The water-stream demo (a charged rod bending a thin stream of water) is popular, but be careful how you explain it. It is often presented as dipoles rotating to face the rod, but induced charges on the stream play a large part, and non-polar liquids can be deflected too. If you use it, present it as a puzzle rather than as proof of polarity.

Differentiation

Support

• A writing frame: “The intermolecular forces in ___ are ___. They are stronger/weaker than in ___ because ___. So more/less energy is needed to ___.”
• Pre-drawn water molecules with lone pairs marked, so students only add the hydrogen bonds.
• A flowchart for identifying forces: “Does it have H–N, H–O or H–F? → hydrogen bonding. Is it polar? → dipole–dipole. Always → London.”

Stretch

• Explain why ice is less dense than liquid water (the open hydrogen-bonded lattice).
• Compare the boiling points of the noble gases and link them to the number of electrons.
• Predict and explain the solubility of ethanol, hexane and sodium chloride in water, using the idea that a solute dissolves when new attractions can replace the ones broken.
• Use the van der Waals calculator to see how the attraction constant a is larger for gases with stronger intermolecular forces.

Question bank

Recall

1. Name the three types of intermolecular force. (London, permanent dipole–dipole, hydrogen bonding.)
2. Which elements must a hydrogen atom be bonded to for hydrogen bonding? (N, O or F.)

Application 3. Which has the higher boiling point, HCl or HBr, and why? (HBr: more electrons, stronger London forces, which outweigh the smaller dipole.) 4. Identify all the intermolecular forces in liquid methanol. (London, dipole–dipole and hydrogen bonding.)

Explanation 5. Explain why butane boils at a higher temperature than propane. (Larger molecule, more electrons and more surface contact: stronger London forces to overcome.) 6. Explain why ammonia has a higher boiling point than phosphine, PH₃, even though PH₃ has more electrons. (Ammonia forms hydrogen bonds; PH₃ cannot, because P is not electronegative enough.)

Challenge 7. Hexane and water don’t mix, but ethanol and water mix in all proportions. Explain in terms of forces broken and formed.

Key takeaways

• Separate “bonds within” from “forces between” in lesson 1, and keep asking which one is overcome.
• Teach London forces as universal, then add dipole–dipole and hydrogen bonding on top.
• Real boiling-point data, especially the hydride graph, gives students evidence rather than rules.
• Low-risk activities (evaporation race, surface tension, model kits) make the forces visible without hazardous chemicals.
• Expect and pre-empt the classic errors: boiling breaking bonds, “any H” hydrogen bonding, and polar bonds equalling polar molecules.