London Forces
A concise revision guide to induced dipole–dipole interactions (London forces): how instantaneous dipoles induce dipoles in neighbouring molecules, why every molecule has them, and how the number of electrons and molecular shape control their strength, for OCR A A Level Chemistry.
GCSE Recap: Counting Electrons
Three quick GCSE questions on electrons before you meet the force that their movement creates.
What London Forces Are
Induced dipole–dipole interactions (London forces) are weak attractions between every molecule and every separate atom, caused by the constant movement of electrons. They are the only intermolecular force that exists between non-polar molecules such as N₂, CH₄ and I₂, and between the atoms of the noble gases.
The electrons in a molecule are never still. At any instant the electron cloud can be slightly denser on one side than the other, so that side is momentarily δ- and the opposite side δ+. This is an instantaneous dipole. It lasts only for a fraction of a second, but while it exists it pushes the electrons of a neighbouring molecule away from the δ- end and pulls them towards the δ+ end, creating an induced dipole in the neighbour. The two dipoles then attract each other.
Because the instantaneous dipole keeps changing, the induced dipole changes with it, so the attraction is always there on average even though no single dipole is permanent. This is why students should describe the force as instantaneous dipole-induced dipole rather than as a fixed charge on the molecule.
An instantaneous dipole in one molecule induces a dipole in its neighbour; the δ- end of one attracts the δ+ end of the other.
Definition: London forces are attractions between an instantaneous dipole in one molecule and the dipole it induces in a neighbouring molecule.
Quick Check: The Mechanism Step by Step
Build the explanation for two ethene molecules, in the order an examiner wants it.
Why Every Molecule Has Them
Any particle with electrons can have an instantaneous dipole, so London forces act between all molecules and atoms, polar or not. A polar molecule such as HCl has London forces as well as the attraction between its permanent dipoles; the two forces add together.
This is the point most often missed in exam answers. When a question asks for the intermolecular forces in a polar molecule, the answer must include London forces alongside the stronger force, because they never switch off.
Common mistake: Writing that “polar molecules have permanent dipole forces instead of London forces”. They have both. The word to use is in addition to.
Quick Check: Instead Of, or In Addition To
Judge one sentence written by a student about a polar molecule.
What Decides Their Strength: Number of Electrons
The more electrons a molecule has, the larger and more polarisable its electron cloud, so the larger the instantaneous and induced dipoles and the stronger the London forces. Molecules with many electrons therefore need more energy to separate and have higher boiling points.
The noble gases show the effect with no other force present. Each is a single atom, so the only attraction between the atoms is the London forces between them, and the boiling point rises steadily as the number of electrons rises.
| Noble gas | Electrons per atom | boiling point / °C |
|---|---|---|
| Helium, He | 2 | -269 |
| Neon, Ne | 10 | -246 |
| Argon, Ar | 18 | -186 |
| Krypton, Kr | 36 | -153 |
| Xenon, Xe | 54 | -108 |
The halogens follow the same pattern. Fluorine and chlorine are gases at room temperature, bromine is a liquid and iodine is a solid, because the number of electrons per molecule rises from 18 in F₂ to 106 in I₂ and the London forces between the molecules strengthen down the group.
| Halogen | Electrons per molecule | melting point / °C | boiling point / °C | State at 25 °C |
|---|---|---|---|---|
| F₂ | 18 | -220 | -188 | gas |
| Cl₂ | 34 | -101 | -34 | gas |
| Br₂ | 70 | -7 | 59 | liquid |
| I₂ | 106 | 114 | 184 | solid |
Boiling point rises with the number of electrons for both series because the London forces between the particles become stronger.
Exam sentence: I₂ has a higher boiling point than Cl₂ because I₂ molecules have more electrons, so the London forces between the molecules are stronger and more energy is needed to overcome them.
Quick Check: Which of the Pair Boils Higher
Four quick comparisons using species that are not on this page, including one beyond the end of the noble gas table.
What Decides Their Strength: Shape and Surface Contact
Molecules with the same number of electrons can still differ, because the attraction depends on how closely the electron clouds of neighbouring molecules can approach. Long, straight molecules can lie alongside each other with a large area of contact, so their London forces are stronger than those between compact, branched molecules of the same formula.
Pentane and 2,2-dimethylpropane both have the formula C₅H₁₂ and 42 electrons. Pentane is a straight chain and boils at 36 °C; 2,2-dimethylpropane is almost spherical and boils at 10 °C. The branched molecule cannot get as close to its neighbours, so less energy is needed to separate them. The attraction between two molecules is really the sum of many small attractions between the parts of their surfaces that lie close together, so more contact simply means more of them to overcome. The Boiling Temperature Trends page uses this idea for the whole alkane series.
Key idea: Two factors set the size of London forces: the number of electrons (the larger factor) and the surface area of contact between neighbouring molecules.
Quick Check: Same Electrons, Different Shape
Two isomers with identical electron counts boil 19 °C apart. Decide why.
How Weak Is Weak?
London forces are the weakest of the three intermolecular forces for molecules of a similar size, but “weak” is relative to covalent bonds, not to nothing. Roughly, a covalent bond needs 150 to 1000 kJ mol⁻¹ to break, a hydrogen bond 10 to 40 kJ mol⁻¹, and the London forces between two small molecules only a few kJ mol⁻¹. For a very large molecule such as a polymer chain or I₂, however, the London forces add up along the whole molecule and can exceed a hydrogen bond in total. Iodine shows it: nothing else holds one I₂ molecule to the next, yet about 62 kJ mol⁻¹ is needed to turn one mole of the solid straight into gas, more than any single hydrogen bond.
| Interaction | Typical energy to break / kJ mol⁻¹ | Acts between |
|---|---|---|
| Covalent bond | 150 to 1000 | atoms within a molecule |
| Ionic bond (lattice) | 600 to 4000 | ions; this is the energy needed to separate one mole of the lattice into gaseous ions |
| Metallic bond | 80 to 900 | metal ions and delocalised electrons |
| Hydrogen bond | 10 to 40 | molecules with N-H, O-H or H-F |
| permanent dipole–dipole interactions | 3 to 20 | polar molecules |
| London forces | 1 to 20, rising with size | all molecules and atoms |
This ladder explains why simple molecular substances melt and boil at low temperatures: only the intermolecular forces are overcome, never the covalent bonds inside the molecules. When ice melts or water boils, every O-H bond stays intact.
Exam trap: Never write that covalent bonds break when a molecular substance boils. The molecules separate; the bonds within them do not.
Names You Will See
OCR names the weakest interaction induced dipole–dipole interactions and allows London forces as an alternative; both permanent and induced dipole–dipole interactions can be called van der Waals’ forces.
Whatever name the question uses, the explanation is the same: moving electrons, an instantaneous dipole, an induced dipole in the neighbour, and an attraction whose size grows with the number of electrons and the area of contact.
Exam wording: Use the name your specification uses: “London forces”. Then say what causes them, because the name alone earns no marks.
Quick Check: The Ideas That Catch People Out
Pick the one accurate statement in each round to finish the page.
Common Exam Points
Explain why boiling point rises down Group 17
More electrons per molecule, larger instantaneous and induced dipoles, stronger London forces, more energy needed to separate the molecules.
Explain why a branched isomer boils lower
Same number of electrons, but a smaller surface area of contact between molecules, so weaker London forces.
State the forces between molecules of a polar substance
Permanent dipole–dipole interactions in addition to London forces; if the molecule has N-H, O-H or H-F, hydrogen bonds as well.
Do not say
“The bonds break” for boiling; “London forces only occur in non-polar molecules”; “the dipole is permanent”.
FAQs
Use these quick answers to check the London forces ideas that appear most often in OCR A A Level Chemistry questions.
What are London forces?
They are attractions between an instantaneous dipole in one molecule and the dipole it induces in a neighbouring molecule. The electrons in any particle are constantly moving, so at any instant the electron cloud can be uneven, and that temporary dipole induces an opposite dipole in the neighbour.
Do London forces act between all molecules?
Yes. Every molecule and every separate atom has electrons, so every one of them has London forces. A polar molecule has them in addition to its permanent dipole–dipole interactions, never instead of them.
What makes them stronger?
Two things. The number of electrons is the larger factor: more electrons mean a bigger, more polarisable cloud and larger induced dipoles. The shape matters too, because long straight molecules touch their neighbours over a larger area than compact branched ones.
Why does boiling point increase down Group 17?
The number of electrons per molecule rises from 18 in F₂ to 106 in I₂, so the London forces between the molecules get stronger and more energy is needed to separate them. That is why iodine is a solid and chlorine a gas at room temperature.
Are van der Waals forces the same thing?
OCR names the weakest interaction induced dipole–dipole interactions and allows London forces as an alternative; both permanent and induced dipole–dipole interactions can be called van der Waals’ forces.
Copyright and author footprint: This OLS revision page was written for Online Learning System by Dr. Mohammed Al-Fatah. It is designed for A Level Chemistry revision and should not be copied or redistributed without permission.
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