Why does helium ever become a liquid?
Helium atoms do not form bonds. Each atom has a full outer shell, an even, symmetrical electron cloud and no permanent dipole. There seems to be nothing to make one helium atom attract another. And yet if you cool helium to about 4.2 K, which is −269 °C, it condenses into a liquid. Liquid means the atoms are clinging together. Something must be pulling them.
That something is the weakest, and yet the most universal, force between particles: the London dispersion force. The polar molecules lecture ended by saying that the next topic would look at how molecules hold on to each other. It starts here, because London forces act between every pair of molecules in existence, polar or not.
Forces inside and forces between
First, a distinction that the whole topic depends on. A covalent bond holds atoms together inside a molecule. It is strong: breaking the H–H bond takes 436 kJ/mol. An intermolecular force is an attraction between separate molecules. These are much weaker, and they are what you overcome when a liquid boils or a solid melts.
When water boils, what is inside the bubbles?
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Water molecules, H₂O, as a gas. Not hydrogen and oxygen. Boiling separates whole molecules from each other by overcoming the intermolecular forces. The covalent O–H bonds inside each molecule stay intact. If boiling broke covalent bonds, a kettle would release hydrogen gas. So whenever you compare boiling points, you are comparing the forces between molecules.
A dipole that lasts an instant
Electrons are not fixed in place. On average, a helium atom's electron cloud is perfectly even around the nucleus. But at any single instant, the electrons may happen to be a little more on one side than the other. For that instant, the atom has a slightly negative side (δ−) and a slightly positive side (δ+). This is called an instantaneous dipole. A moment later it has shifted, faded or reversed.
Now bring a second atom close. The δ− side of the first atom repels the electrons of the second atom and pushes them away. So the second atom's cloud also becomes lopsided, with its δ+ side facing the first atom's δ− side. This is an induced dipole: it was caused by its neighbour. A δ− next to a δ+ means attraction.
Here is the clever part. If A's electrons drift the other way, A's side facing B becomes δ+. That pulls B's electrons towards A, so B's facing side becomes δ−. Again the facing sides are opposite, and again they attract. Whichever way the first dipole points, the dipole it induces lines up to attract it. So although each dipole lasts only an instant and keeps changing, the overall effect is always an attraction. The attraction between an instantaneous dipole and the dipole it induces is the London dispersion force.
Watch it happen below.
Choose an atom. Press Play to watch the electron clouds flicker, or press Next instant to step. Watch the δ signs on the facing sides, and the attraction line. Compare a small atom with a large one.
What the simulation shows, in words:
- Any instant: if A's electrons drift towards B, A's facing side is δ− and B's facing side becomes δ+. If they drift away from B, A's facing side is δ+ and B's becomes δ−. Either way the facing sides attract.
- Small atom (helium, 2 electrons): a small, tightly held cloud. The shifts are tiny, so the attraction is very weak. Helium boils at about 4.2 K.
- Large atom (xenon, 54 electrons): a large cloud whose outer electrons are far from the nucleus. The shifts are much bigger, so the attraction is much stronger. Xenon boils at about 166 K.
London forces are named after the physicist Fritz London, who explained them using quantum mechanics around 1930. Some books call them dispersion forces, or include them under the wider name van der Waals forces. This course says "London forces" after this first mention.
A tug of war: attraction against motion
If London forces act between all particles at all times, why is argon a gas at room temperature? Because attraction is only half the story. The particles of any substance are always moving, and the higher the temperature, the faster they move. Each pair of neighbours is in a tug of war. The attraction tries to hold them together; their motion tries to carry them apart.
When the particles are moving fast enough to break away from their neighbours, the substance is a gas. Cool it, and the particles slow down. At some temperature, the attraction wins often enough for the particles to cling together, and the gas condenses into a liquid. That temperature is the boiling point. So a boiling point measures how strong the attraction is: the stronger the forces between particles, the more motion it takes to pull them apart, and the higher the boiling point.
Argon is a gas at room temperature. Does that mean there are no London forces between argon atoms at room temperature?
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No. The forces are still there. At room temperature, about 298 K, the argon atoms move so fast that the weak London forces cannot hold them together, so argon is a gas. Cool it below 87.3 K and the atoms slow down enough for the same forces to win, and argon becomes a liquid. The force did not switch on at 87.3 K; the motion became too weak to overcome it.
What makes London forces stronger
The size of a London force depends on how easily the electron cloud can be pushed out of shape. Chemists call this polarisability. A more polarisable cloud makes bigger instantaneous dipoles and responds more strongly to its neighbours. Two things increase it.
- More electrons. Bigger atoms and molecules have more electrons, and their outer electrons are further from the nucleus and more shielded from it, so they are held less tightly and shift more easily.
- More contact. Molecules attract only where they come close. Long, thin molecules can lie alongside each other and touch along their whole length. Compact, round molecules touch only at a few points.
Why should more electrons make a cloud easier to push around? The periodic trends lectures gave the reason. Going down a group, each atom has an extra shell. Its outer electrons are further from the nucleus and more shielded by the inner shells, so they feel a weaker pull. That is why ionisation energy falls down a group: the outer electrons are easier to remove. The same weak grip makes them easier to shift to one side. So the trend that makes xenon easier to ionise than helium also makes it more polarisable.
The cleanest evidence comes from the noble gases. They are single atoms with no bonds and no permanent dipoles, so London forces are the only attraction between them. Going down the group, only the number of electrons changes.
| Noble gas | Electrons | Boiling point (K) | Boiling point (°C) |
|---|---|---|---|
| Helium | 2 | 4.2 | −269 |
| Neon | 10 | 27.1 | −246 |
| Argon | 18 | 87.3 | −186 |
| Krypton | 36 | 120.9 | −152 |
| Xenon | 54 | 166.1 | −107 |
Boiling points from Chemguide; other data tables differ by about 1 K. Every step down the group adds electrons, makes the London forces stronger and raises the boiling point. The pattern is smooth because nothing else changes.
Worked examples
F₂ and I₂ are both non-polar molecules made of two identical atoms, so only London forces act between their molecules. Explain why fluorine is a gas at room temperature and iodine is a solid.
F₂ has electrons. I₂ has . Iodine's large electron cloud is far more polarisable, so the London forces between I₂ molecules are much stronger. Fluorine boils at 85 K (−188 °C). Iodine boils at 457 K (184 °C), so at room temperature its molecules are held firmly enough to form a solid. Same kind of molecule, same kind of force, but a difference of more than 370 degrees in boiling point.
Pentane, C₅H₁₂, can be built in three ways: a straight chain (pentane), a chain with one branch (methylbutane, often called isopentane), and a compact ball with four CH₃ groups around a central carbon (dimethylpropane, often called neopentane). All three have the same formula, so all three have exactly 42 electrons. Yet they boil at 36 °C, 27 °C and 9.5 °C.
The difference is contact. Straight-chain pentane molecules can lie side by side and touch along their whole length, so there are many points of attraction. The ball-shaped neopentane molecules touch only at a small area, so the London forces between them are weaker. Same number of electrons, different shape, different strength.
Both effects work together along the family of straight-chain alkanes. Each extra CH₂ group adds 8 electrons and makes the chain longer, so it adds more electrons and more contact. The smallest members, such as methane in natural gas, are gases at room temperature. Middle-sized ones, like those in petrol, are liquids. Very long chains, like those in candle wax, are solids. One family, one kind of force, three states of matter, decided by chain length.
Before you read on, make a prediction about molecules with the same formula.
Hexane, C₆H₁₄, can also be built as a straight chain or as a compact, highly branched molecule (2,2-dimethylbutane). Both have the same formula and the same mass. Which has the higher boiling point?
The straight-chain hexane has the higher boiling point. Both molecules have the same 50 electrons, but the straight chains can lie alongside each other and touch along their length, so the London forces between them are stronger. The compact branched molecule touches its neighbours over a smaller area. This is the same pattern as the three pentanes.
Do London forces act between water molecules, which are polar?
Yes. Every molecule has electrons, and every electron cloud flickers. So London forces act between all molecules, polar or non-polar. In polar molecules, other attractions act as well, and the next two lectures look at them. In small polar molecules like water, those other forces are larger than the London forces. In big molecules, London forces can be the largest attraction of all, even when the molecules are polar.
Three classic mistakes
They act between all molecules and atoms. In non-polar substances they are the only intermolecular force, which is why the textbook examples are non-polar. But polar molecules have them too.
Boiling separates whole molecules. The covalent bonds inside them are untouched. Boiling points measure the forces between molecules, not the bonds inside them.
Heavier molecules usually have more electrons, which is what strengthens London forces. But mass is not the cause. The three pentanes have the same mass and boil at very different temperatures. Explain with electrons and contact, not with weight.
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The halogens F₂, Cl₂, Br₂ and I₂ boil at 85 K, 238 K, 332 K and 457 K. Explain the trend, and say which of them are gases, liquids and solids at room temperature (about 298 K).
Hint 1
Count the electrons in each molecule: F has 9, Cl 17, Br 35, I 53.
Hint 2
A substance is a gas at room temperature if its boiling point is below 298 K.
Full answer
The molecules have 18, 34, 70 and 106 electrons. More electrons make a more polarisable cloud and stronger London forces, so the boiling point rises down the group. F₂ (85 K) and Cl₂ (238 K) boil below 298 K, so they are gases. Br₂ boils at 332 K, above room temperature, and it is a liquid. I₂ boils at 457 K and is a solid at room temperature.
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Isopentane (one branch) boils at 27 °C, between pentane (36 °C) and neopentane (9.5 °C). Explain why.
Hint 1
All three have 42 electrons.
Hint 2
Compare how much each shape lets molecules touch.
Full answer
With the same number of electrons, the difference is contact area. One branch makes isopentane more compact than the straight chain but less compact than the ball-shaped neopentane. So its molecules touch over an area between the other two, its London forces are in between, and so is its boiling point.
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A student says: "Methane has almost no bond polarity and no dipole, so nothing attracts its molecules." Methane does condense to a liquid when cooled enough. Use that fact to explain what is wrong.
Hint 1
If there were no attraction at all, could a gas ever become a liquid?
Hint 2
Methane has 10 electrons that can form instantaneous dipoles.
Full answer
Any substance that can be liquefied must have some attraction between its particles, or they would never cling together. Methane's 10 electrons form instantaneous dipoles that induce dipoles in neighbouring molecules, so London forces act. They are weak, which is why methane has to be cooled a long way before it condenses, but they are not zero.
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Geckos can cling to smooth glass. A study by Autumn and colleagues (2002) found evidence that the many microscopic hairs on their feet stick by van der Waals forces. Use the ideas of this lecture to explain why a gecko's foot sticks but a smooth rubber ball pressed against glass does not.
Hint 1
London forces act only where surfaces are extremely close, a few atoms apart.
Hint 2
Which surface can get extremely close to the glass at many more points?
Full answer
Each London attraction is tiny and works only over very short distances. A ball, however smooth it looks, touches the glass at only a few points at that tiny distance, so the total force is small. A gecko's foot is covered in a huge number of very fine, flexible hairs that bend to follow the glass and bring an enormous number of tiny contact points close enough. Many weak forces add up to a large one. It is the pentane lesson on a bigger scale: contact area decides the strength.
One London force is tiny, far weaker than a covalent bond. But they add up. A long molecule touches its neighbours at many points, and the total can be large. That is why candle wax, made of long non-polar molecules, is a solid at room temperature, and why iodine is a solid while fluorine is a gas.
London forces are the background attraction present between all molecules. Polar molecules add a second, permanent attraction on top: their δ+ and δ− ends pull on each other all the time, not just for an instant. That is the next lecture.
- Intermolecular forces act between molecules; boiling and melting overcome them, not the covalent bonds inside molecules.
- An instantaneous dipole in one particle induces a dipole in its neighbour, and the two always attract: this is the London dispersion force.
- London forces act between all atoms and molecules, polar or non-polar.
- They are stronger with more electrons (a more polarisable cloud) and with more contact between molecules.
- Evidence: noble gas and halogen boiling points rise with electron count; straight-chain pentane boils higher than compact neopentane.