A puzzle from two lectures ago
The polar molecules lecture showed that chloroform, CHCl₃, is polar, and that tetrachloromethane, CCl₄, is not: its four C–Cl dipoles cancel. So you might expect chloroform molecules to attract each other more strongly, and chloroform to boil at the higher temperature. It does not. CCl₄ boils at 77 °C and CHCl₃ at 61 °C. The non-polar liquid boils 16 degrees higher.
The London forces lecture ended by promising a second attraction, one that polar molecules have on top of London forces. This lecture introduces it, the permanent dipole–dipole force, and then solves the puzzle. The answer turns on a habit every chemist needs: compare like with like.
Dipoles that stay
A London force comes from a dipole that flickers in and out of existence. A polar molecule is different. Hydrogen chloride, H–Cl, always has a δ+ hydrogen end and a δ− chlorine end, because chlorine is more electronegative (the difference is ). That dipole is permanent.
When polar molecules are close together, they tend to turn so that the δ+ end of one faces the δ− end of a neighbour. Opposite charges attract, so this arrangement holds them together. Like charges facing each other would repel, so that arrangement is avoided.
Polar molecules attract each other because the δ+ end of one is drawn to the δ− end of another. These forces act in addition to the London forces that act between all molecules.
In a liquid, molecules are not neatly lined up like the figure. They tumble and move all the time. But at any moment, the attracting arrangements are a little more common than the repelling ones, because molecules that turn the right way are pulled together. So on average there is a net attraction. Heat makes the molecules move and turn faster, which scrambles the lining-up. Explore this below.
Each arrow is a polar molecule, pointing from δ+ to δ−. Choose a temperature, then press Shake to see another instant. Compare how well the molecules line up head to tail.
What the explorer shows, in words:
- Cold: the molecules move slowly, so most of them stay lined up head to tail, δ+ next to δ−. The average attraction is strong.
- Warm: the molecules turn more, so the lining-up is only partial. There is still a clear net attraction.
- Hot: the molecules point in almost random directions. Attractions and repulsions nearly cancel, so the net dipole–dipole attraction is weak.
This temperature effect is one difference between the two kinds of force. London forces come from electrons flickering inside each molecule, which does not depend much on how the molecules are turned. Dipole–dipole attraction depends on the molecules staying lined up, so it weakens as the temperature rises.
Think about the three states of matter in these terms. In a gas, molecules are far apart and move fast, so they hardly feel each other's dipoles at all. In a liquid, they are close together but tumbling, so they line up only partly, and only on average. In a solid, the molecules are held in place, so polar molecules can settle into arrangements where most neighbours face each other the right way. So the colder a polar substance is, the more fully its dipoles can line up and pull on each other.
Which molecules have dipole–dipole forces?
Only molecules with an overall dipole. Polar bonds are not enough. The polar molecules lecture showed that a molecule can have polar bonds whose dipoles cancel, as in CO₂, BF₃ and CCl₄. Such a molecule has no δ+ end and no δ− end as a whole, so it has nothing to line up. Between those molecules, the only attraction is London forces.
Each C=O bond in carbon dioxide is polar (difference 0.89). Are there dipole–dipole forces between CO₂ molecules?
Show the answer
No. CO₂ is linear, so its two bond dipoles cancel and the molecule has no overall dipole. A CO₂ molecule has slightly negative oxygens at both ends and a slightly positive carbon in the middle, but no positive end and negative end to line up with a neighbour's. So the attraction between CO₂ molecules is London forces only. To decide whether dipole–dipole forces act, use the shape-based test from the polar molecules lecture, not the bonds alone.
The size of the molecular dipole matters too. The bond polarity lecture measured HCl's dipole moment at about 1.11 debye. A molecule with a larger dipole moment has bigger partial charges or more widely separated ones, so it attracts its neighbours more strongly. Going from HCl to HI, the electronegativity difference falls from 0.96 to 0.46, the dipole gets smaller, and the dipole–dipole part of the attraction gets weaker. Keep that in mind for the hydrogen halides below.
Fair tests: same electrons, different polarity
Every polar molecule has London forces too. So to see what the permanent dipole adds, you need a fair test. Compare a polar and a non-polar substance with about the same number of electrons, so that their London forces are about the same.
Hydrogen chloride has electrons. An argon atom also has 18. So their London forces should be similar. But HCl is polar and argon is not.
HCl boils at about −85 °C; argon at about −186 °C. The polar substance boils about 100 degrees higher. With the London forces matched, that difference comes from the dipole–dipole forces that HCl has and argon lacks.
Propanone (acetone), CH₃COCH₃, has 32 electrons and a polar C=O group (difference 0.89). 2-Methylpropane, C₄H₁₀, has 34 electrons and is non-polar. Both molecules are compact and of similar size.
Propanone boils at 56 °C; 2-methylpropane at −12 °C. A difference of 68 degrees, even though the non-polar molecule has slightly more electrons. Again, the permanent dipoles make the difference. (Propanone cannot form hydrogen bonds with itself, because it has no O–H group. So this really is a dipole–dipole effect.)
When London forces win
Now look at a series where polarity and size change together: the hydrogen halides.
| Molecule | Electronegativity difference | Electrons | Boiling point |
|---|---|---|---|
| HCl | 0.96 | 18 | about −85 °C |
| HBr | 0.76 | 36 | about −67 °C |
| HI | 0.46 | 54 | about −35 °C |
Going down the table, the molecules become less polar, so their dipole–dipole forces get weaker. Yet the boiling points rise. The reason is the electron count, which triples from HCl to HI. The London forces grow much faster than the dipole–dipole forces shrink. So for this trend, London forces are the deciding factor, even though all three molecules are polar.
Hydrogen fluoride is missing from the table on purpose. HF boils at about 20 °C, far higher than the trend would suggest. That is because of hydrogen bonding, a much stronger special case, and it is the subject of the next lecture.
Now the opening puzzle. CHCl₃ has 58 electrons; CCl₄ has 74. CCl₄'s extra chlorine atom brings 16 more electrons and a larger, more polarisable surface. Its stronger London forces outweigh the dipole–dipole forces that chloroform has and CCl₄ lacks. So CCl₄ boils higher, at 77 °C against 61 °C. Polarity does matter, but it is not the only thing that matters.
A student says: "The more polar a molecule is, the higher its boiling point." Is that a safe rule?
Show the answer
Only with a condition: when the molecules have about the same number of electrons and a similar shape. HCl boils higher than argon, and propanone higher than 2-methylpropane, because those pairs are matched. But HCl is more polar than HI and boils lower, and CHCl₃ is more polar than CCl₄ and boils lower, because in those pairs the bigger molecule has much stronger London forces. A safe rule names both factors: electron count (and shape) for London forces, and polarity for dipole–dipole forces.
Sort these comparisons by the main reason for the difference.
For each pair, decide what mainly explains why the first boils higher. Check the electron counts first: if they are about equal, look at polarity.
Eight comparisons sorted, each with its reason.
- Mainly London forces: HI above HCl; HBr above HCl; CCl₄ above CHCl₃; I₂ above F₂; straight-chain pentane above neopentane.
- Mainly dipole–dipole forces: HCl above argon; propanone above 2-methylpropane; HCl above F₂.
Putting the forces in order
You now know two kinds of force between molecules, and the next lecture adds a third. How do they compare? For molecules of about the same size, dipole–dipole forces add noticeably to London forces, and hydrogen bonds add much more. But size can overturn that order: the London forces between large molecules can be stronger than all the other attractions between small ones. Iodine, with only London forces, is a solid at room temperature, while hydrogen chloride, with London and dipole–dipole forces, is a gas. That is why every comparison should start with the same question: are these molecules about the same size? If they are, look at polarity. If they are not, look at the electrons first.
Microwave ovens: a popular myth
You may have heard that microwave ovens work because their frequency "matches the natural frequency of water molecules". That is a myth. Christopher Baird, a physicist at West Texas A&M University, points out that ovens work around 2.45 GHz mainly because that band was set aside for them. Water has no special resonance there. Its first resonance is at a far higher frequency.
What really happens connects to this lecture. The oven's electric field reverses direction billions of times a second. Polar molecules, such as water, keep trying to turn and line up with it, and as they turn they bump into their neighbours. That jostling turns the wave's energy into heat. Many polar molecules in food absorb energy this way, not just water. So the permanent dipole matters, but there is no magic matching frequency.
Three classic mistakes
They have both. London forces act between all molecules. Dipole–dipole forces are an extra attraction that polar molecules have on top.
Only when the electron counts and shapes are similar. HCl is more polar than HI but boils lower; CHCl₃ is polar and CCl₄ is not, yet CCl₄ boils higher.
HF, H₂O and NH₃ have far higher boiling points than their size suggests. That is hydrogen bonding, a special and much stronger case. Keep it separate; the next lecture explains it.
Use Pauling electronegativities where you need them: H 2.20, C 2.55, O 3.44, F 3.98, Cl 3.16, Br 2.96, I 2.66.
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HCl boils at about −85 °C and F₂ at about −188 °C. Both molecules have 18 electrons. Explain the difference.
Hint 1
Same electron count means similar London forces.
Hint 2
Which of the two molecules is polar?
Full answer
With 18 electrons each, their London forces are similar. F₂ is made of two identical atoms, so its bond is non-polar. HCl is polar (difference 0.96), so it also has permanent dipole–dipole forces. The extra attraction raises its boiling point by about 100 degrees.
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Work out the electronegativity difference for HCl, HBr and HI. Use your answers, and the boiling points in the table, to explain why the trend cannot be due to dipole–dipole forces.
Hint 1
Subtract hydrogen's 2.20 from each halogen's value.
Hint 2
As the difference falls, what should happen to dipole–dipole forces? What actually happens to the boiling point?
Full answer
HCl 0.96, HBr 0.76, HI 0.46. Polarity falls down the group, so dipole–dipole forces get weaker. If they decided the trend, the boiling points would fall. Instead they rise, from about −85 °C to about −35 °C. So something else is growing: the London forces, as the electron count rises from 18 to 54.
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CCl₄ (non-polar) boils at 77 °C and CHCl₃ (polar) at 61 °C. Explain this without saying "CCl₄ is heavier".
Hint 1
Count the electrons: C has 6, H has 1, Cl has 17.
Hint 2
Which force depends on the number of electrons?
Full answer
CCl₄ has electrons; CHCl₃ has . CCl₄'s larger, more polarisable electron cloud gives stronger London forces. Chloroform has dipole–dipole forces as well, but they are smaller than the extra London forces of CCl₄. So the London forces decide.
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Explain why dipole–dipole attraction between polar molecules gets weaker, on average, as the temperature rises, while London forces are much less affected.
Hint 1
What do polar molecules have to do for their dipoles to attract?
Hint 2
What does a higher temperature do to the motion of molecules?
Full answer
Dipole–dipole attraction needs neighbouring molecules to be turned the right way, δ+ towards δ−. At higher temperatures, molecules move and rotate faster, so they spend less time lined up and more time in random or repelling orientations. The net attraction falls. London forces come from fluctuations in each molecule's own electron cloud. These induce dipoles in neighbours whichever way the molecules are turned, so they depend much less on how the molecules are lined up.
Every boiling point is a sum of attractions. London forces are always there, and they grow with the number of electrons and the contact between molecules. Polar molecules add permanent dipole–dipole forces on top. When you compare two substances, first ask which forces each has, then compare the sizes: electron counts for London forces, electronegativity differences and shape for dipoles. Most surprises, like CCl₄ boiling above CHCl₃, disappear once you count both.
One case does not fit: HF boils far above HCl, even though it has fewer electrons. Water and ammonia behave the same way. The next lecture explains this third, much stronger force between molecules: hydrogen bonding.
- Polar molecules attract each other through permanent dipole–dipole forces: δ+ of one towards δ− of another.
- These forces act in addition to London forces, which every molecule has.
- With similar electron counts, a polar substance boils higher than a non-polar one (HCl above Ar; propanone above 2-methylpropane).
- When electron counts differ a lot, London forces usually decide (HCl < HBr < HI; CCl₄ above CHCl₃).
- Higher temperatures scramble the lining-up of dipoles, so dipole–dipole attraction weakens on average.