Intermolecular Forces: London, Dipole-Dipole and Hydrogen Bonding
Intermolecular forces are the attractions between whole molecules, not the bonds inside them. Get that distinction straight and the rest of the chapter follows. Get it wrong and nothing makes sense, which is why most learners find this topic harder than it needs to be.
This page covers the three types you need for Grade 11, what Van der Waals actually means, and the prescribed practical that measures all five properties these forces control.
The mistake that ruins this chapter
| What it is | Broken when you | |
|---|---|---|
| Intramolecular | The bonds inside a molecule. Covalent, ionic, metallic | React the substance chemically |
| Intermolecular | The attractions between molecules | Melt or boil it |
When you boil water you do not break the O to H bonds. If you did, you would get hydrogen and oxygen gas, and steam is not that. You separate whole water molecules from each other, and the things holding them together are intermolecular forces.
Intermolecular forces are much weaker than the bonds inside the molecule. That is why boiling a kettle takes a few minutes and splitting water into its elements takes an electrolysis rig.
The three you need to know
London forces
Present in every single substance, and the only force between non-polar molecules.
Electrons are always moving. At any instant they can sit slightly more to one side of a molecule than the other, which makes a brief, tiny dipole. That flicker induces a matching dipole in the neighbour, and the two attract.
They are the weakest of the three, but they grow with the size of the molecule. A bigger molecule has a bigger, floppier electron cloud that distorts more easily, and more surface for the neighbour to feel. This is why iodine is a solid and chlorine is a gas, even though both are non-polar diatomic molecules with nothing else going on.
Also called dispersion forces, or induced-dipole induced-dipole forces. Same thing, three names.
Dipole-dipole forces
Between polar molecules.
When two bonded atoms differ in electronegativity, the shared electrons sit closer to one of them. That end goes slightly negative, the other slightly positive, and the molecule is polar. The positive end of one molecule then attracts the negative end of the next.
Stronger than London forces, because the dipole is permanent rather than a flicker.
Hydrogen bonding
The strongest of the three, and it has strict entry requirements.
You only get hydrogen bonding when a hydrogen atom is bonded directly to nitrogen, oxygen or fluorine. Nothing else counts.
Those three are small and very electronegative, so they strip the hydrogen almost bare, leaving an exposed proton that latches hard onto a lone pair on the next molecule.
This is why water behaves the way it does. Water is a tiny molecule that should boil far below room temperature on size alone. It boils at 100 °C because every molecule can hold several hydrogen bonds at once.
So what is a Van der Waals force?
It is an umbrella term, not a fourth type. And there are two conventions, which is why you will find contradictory answers online.
| Convention | Van der Waals covers | Where you meet it |
|---|---|---|
| Your Grade 11 textbook | London forces, dipole-induced-dipole, dipole-dipole and hydrogen bonding | The flow chart in Chapter 4 |
| Strict physical chemistry | London forces, dipole-induced-dipole and dipole-dipole. Hydrogen bonding sits outside | University, and most of the internet |
In an exam, follow your own textbook. The CAPS Grade 11 book puts hydrogen bonding inside Van der Waals, so that is the answer that earns the mark.
The reason the stricter version exists is that hydrogen bonding is not purely an electrostatic attraction between dipoles. It is directional, it is far stronger than the others, and it has a small amount of covalent character. Chemists beyond school level usually keep it separate for those reasons.
What both conventions agree on: Van der Waals is not a synonym for London forces. London forces are one kind of Van der Waals force. If a question asks you to name the force between two chlorine molecules, London forces is the precise answer and Van der Waals is the vague one.
All six, strongest to weakest
This ordering is asked directly in the Chapter 4 exercises, so it is worth knowing as a list.
| Force | Example | |
|---|---|---|
| 1 | Ion-dipole | NaCl dissolved in water |
| 2 | Hydrogen bonding | Water, ammonia, HF, ethanol |
| 3 | Ion-induced-dipole | NaCl in CCl4 |
| 4 | Dipole-dipole | ICl, HBr |
| 5 | Dipole-induced-dipole | I2 in water |
| 6 | London forces | Cl2, CH4, helium |
The two ion forces only appear when there are ions present, which in this chapter means a dissolved salt. For a question about a pure liquid or gas, you are choosing between the bottom four.
Which force is it? Work through it in order
- Are there ions involved? If yes, you are looking at ion-dipole (salt in water) or ion-induced-dipole forces
- Is hydrogen bonded to N, O or F? If yes, hydrogen bonding
- Are the molecules polar? If yes, dipole-dipole
- Neither? Then London forces, and they are all you have
London forces are present in all four cases. The question is always which is the strongest force present, because that is the one that decides the physical properties.
| Substance | Strongest force present |
|---|---|
| Cl2, I2, CH4, He | London forces |
| HCl, ICl, HBr | Dipole-dipole |
| H2O, NH3, HF, ethanol | Hydrogen bonding |
| NaCl in water | Ion-dipole |
Carbon dioxide catches people out. Each C to O bond is polar, but the molecule is linear and symmetrical, so the two pull in opposite directions and cancel. CO2 is non-polar and only has London forces.
Where this fits in the curriculum
| Subject | Physical Sciences |
|---|---|
| Grade | 11 |
| Term | 2 |
| Topic | Matter and materials |
| Status | Prescribed formal assessment |
| Marks | 50 |
The experiment: five measurements, one answer
The prescribed practical takes the same everyday liquids and measures five different things about them. Every result comes back to the same cause.
Apparatus
| Item | Qty |
|---|---|
| Beaker, 250 ml | 2 |
| Thermometer, -10 to 110 °C | 1 |
| Test tubes, 25 x 150 mm, rack and holder | 6 |
| Measuring cylinder, 10 ml | 1 |
| Petri dish | 2 |
| Droppers | 1 pk |
| Capillary tubes, about 1 mm bore | 10 |
| Tripod, gauze and a burner | 1 |
You supply the liquids. Nail polish remover, methylated spirits, olive oil, glycerine and table salt are all supermarket or pharmacy items. Four of them are flammable, which makes them dangerous goods in transit, and all of them are cheaper and fresher bought locally on the morning.
Part 1: Evaporation
One drop of each liquid on a glass plate. Time how long it takes to disappear.
| Liquid | Roughly | Why |
|---|---|---|
| Nail polish remover | under 1 min | Acetone. Dipole-dipole only |
| Methylated spirits | 2 to 4 min | Mostly ethanol. Hydrogen bonding |
| Ethanol | 2 to 4 min | Same |
| Water | 15 to 30 min | Hydrogen bonding, small molecule, several bonds each |
Start Part 1 first and walk away from it. This is the single most useful piece of advice on this page. Water takes twenty minutes or more, and a class that stands watching it loses the period. Put the drops down, start the clock, go and do Parts 2 and 3, come back.
And if the water has not gone by the bell, that is your result. Write down "not evaporated after 25 minutes". It is the strongest finding in Part 1, not a failure.
Part 2: Surface tension
Count how many drops you can pile onto a 50c coin before it runs off.
| Liquid | Roughly | Surface tension |
|---|---|---|
| Water | 30 to 35 drops | 72 mN/m |
| Glycerine | 25 to 30 | 63 mN/m |
| Olive oil | 15 to 20 | 32 mN/m |
| Methylated spirits | 10 to 12 | 22 mN/m |
| Nail polish remover | 8 to 12 | 24 mN/m |
Your counts will not match those, and it does not matter. The number depends entirely on how big a drop your dropper makes. The order is the result.
Methylated spirits and nail polish remover are within 2 mN/m of each other, so do not expect a reliable difference between those two. Everything else separates cleanly.
Two things are happening. Water molecules pull on each other hard, so the surface behaves like a skin and holds a tall dome. And water does not wet the coin well, so it beads instead of spreading. Methylated spirits flattens out the moment it lands and runs straight off the edge.
Part 3: Solubility
Three solutes, three solvents, nine test tubes.
| Water | Ethanol | Olive oil | |
|---|---|---|---|
| Potassium permanganate | Deep purple | Slight purple | No |
| Iodine | Barely, very pale | Brown | Brown |
| Table salt | Dissolves | Barely | No |
Like dissolves like, and this is where that phrase earns its keep.
Salt is ionic. Water is polar, so its partial charges pull the Na⁺ and Cl⁻ ions out of the lattice. Oil has no charges to offer and the salt just sits there.
Iodine is non-polar, held together by weak London forces. It slips easily in among the non-polar oil molecules. To dissolve in water it would have to break hydrogen bonds, and weak London forces are nowhere near a fair trade.
Ethanol is the interesting one. It has an OH end that behaves like water and a hydrocarbon end that behaves like oil, so it dissolves a little of everything and not much of anything.
The chloroform problem, and what to use instead
The textbook uses chloroform as the third solvent. Do not.
Trichloromethane is a suspected human carcinogen with liver and kidney effects on repeated exposure, and it is restricted in school laboratories across most of the world.
Olive oil does the same job and it is already on the bench for Parts 2 and 5. Oil against water is also a comparison every learner already understands. If you want a strictly non-polar alkane, medicinal liquid paraffin from a pharmacy is better still and costs almost nothing.
Part 4: Boiling point
Measure water and ethanol. Look up the rest.
| Liquid | Boiling point | The force responsible |
|---|---|---|
| Nail polish remover | 56 °C | Dipole-dipole only |
| Ethanol | 78 °C | Hydrogen bonding, one OH |
| Methylated spirits | about 78 °C | Mostly ethanol |
| Water | 100 °C | Hydrogen bonding, and a very small molecule |
| Glycerine | 290 °C | Three OH groups |
| Olive oil | about 300 °C | Enormous molecules. London forces add up |
This table is the best thing in the whole experiment, because it shows two effects at once.
Read the top four and you see the type of force mattering: acetone, then ethanol, then water, as the forces get stronger.
Read the bottom two and you see size mattering. Glycerine and olive oil are not more polar than water. Glycerine has three OH groups instead of water's two hydrogens, and olive oil molecules are simply vast, with far more surface for London forces to work over.
Heat the ethanol in a water bath, never over the flame. Ethanol vapour catches fire easily. A water bath cannot exceed 100 °C, which is comfortably above the 78 °C the ethanol needs.
Part 5: Capillarity
Stand a narrow tube in each liquid and measure how far it climbs.
| Liquid | Height risen in a 1 mm tube |
|---|---|
| Water | about 30 mm |
| Glycerine | about 20 mm |
| Olive oil | about 14 mm |
| Methylated spirits | about 12 mm |
| Nail polish remover | about 12 mm |
Glycerine and olive oil are slow. Both are thick and take a minute or two to reach their final height, so do not measure them the instant the tube goes in.
Use a capillary tube, not a straw
The textbook offers a drinking straw as the alternative. It does not work.
How far a liquid climbs is inversely proportional to the width of the tube. A capillary tube is about 1 mm across. A straw is about 5 mm.
| Tube | Water climbs |
|---|---|
| Capillary tube, 1 mm | about 30 mm |
| Drinking straw, 5 mm | about 6 mm |
Six millimetres, when the real differences between your liquids are a few millimetres each. The class cannot separate the results and Part 5 gives them nothing. Capillary tubes cost almost nothing and they are the one item in this practical a school genuinely cannot improvise.
The pattern, and it is the whole point
Put the five results side by side and one liquid wins four of them.
| Evaporation | Surface tension | Boiling point | Capillarity | |
|---|---|---|---|---|
| Water | Slowest | Most drops | Highest measured | Climbs highest |
| Nail polish remover | Fastest | Fewest drops | Lowest | Climbs least |
That is not five facts. It is one fact, measured five ways.
Water molecules hold on to each other through hydrogen bonding, the strongest intermolecular force in the Grade 11 syllabus. Acetone molecules only manage dipole-dipole. Everything in that table follows from those two sentences.
Safety
- Never let potassium permanganate touch glycerine. Poured together they self-heat and burst into flame after about twenty seconds, with no match involved. Both are on the textbook apparatus list, so keep them at opposite ends of the bench and use a separate spatula for each
- Ethanol is heated in a water bath only. A test tube of ethanol over an open flame is a fire
- No chloroform. Use olive oil or liquid paraffin
- Goggles on for the whole practical, not just Part 3
- Nail polish remover and methylated spirits are flammable. Keep both closed and away from the burner, and do not run Parts 1 and 4 on the same bench at once
- Iodine and potassium permanganate stain skin and clothing permanently
If it does not work
| What you see | What caused it |
|---|---|
| The water drop has not evaporated by the bell | Normal, and it is the result. Record how long you waited |
| Every liquid gives about the same drop count | The dropper is running too fast. One drop at a time |
| The coin count changes wildly on the repeat | The coin was not dried between liquids |
| Nothing climbs the capillary tube | The tube is blocked, or still wet from a different liquid. Use a dry tube each time |
| All five climb the same height | You used a straw, or a tube that is too wide |
| The ethanol never reaches 78 °C | The water bath is not hot enough, or the bulb is not in the liquid |
| Iodine gives no colour in water | Correct. Very faint at most, and that is the finding |
| The marking pen ink dissolved | Acetone does that. Mark the glass, not the plastic |
The failure worth keeping
A group that does not dry the coin between liquids will get nonsense, and they should be left to notice it themselves.
A trace of methylated spirits lowers the surface tension of the water that follows and the count collapses. That is contamination changing a measurement, which is the most useful thing anyone takes out of a practical, and it lands far harder when it is their own coin.
How the 50 marks are made up
| Section | Marks |
|---|---|
| Planning, questions and variables | 8 |
| Part 1, evaporation | 7 |
| Part 2, surface tension | 7 |
| Part 3, solubility | 9 |
| Part 4, boiling point | 9 |
| Part 5, capillarity | 6 |
| Conclusion | 4 |
The textbook prints no mark allocation for this one. It says the teacher assesses comprehending, synthesising, generalising, communicating results and drawing conclusions. The split above is ours, weighted towards the two parts that carry the most explaining.
Half the marks are for explanation, not measurement. A group can measure all five things perfectly and score under half if they cannot say why.
The mark most often dropped is not naming the force. "Water has stronger forces" is worth one where "water has hydrogen bonding" is worth two, and that comes up in six separate questions.
If you have time
Put one drop of detergent in the water and redo the coin count. It collapses. Detergent molecules wedge in between the water molecules and break up the hydrogen bonding, which is exactly how soap works.
Add food colouring to the water in Part 5. The rise becomes visible from the back of the room, which is the version to use if you are demonstrating rather than measuring.
Free worksheet and marking memo
Both free, no sign up, straight to the PDF.
- Learner worksheet, 50 marks, with the planning section, all five results tables and the interpretation questions set out
- Marking memorandum, with a worked sample set for all five parts, the mark allocation and a note on the six places learners most often drop marks
Related practicals
- Rate of reaction, Grade 12. Shares the tripod, gauze and burner
- Newton's Second Law, Grade 11. The other prescribed Grade 11 practical
- Ohm's Law, Grade 11. Term 3, and the biggest graph practical of the year
Buy this experiment
We are putting together a complete Intermolecular Forces Kit for Grade 11 with the glassware, thermometer, capillary tubes, heating gear and goggles in one box, plus a printed teacher guide and the marking memo. Coming shortly.
In the meantime every item in the apparatus table above is sold separately. The capillary tubes are the ones to check your store cupboard for. They are the only thing in this practical that cannot be improvised, and a straw will not do.