Permanent Dipole-Dipole Forces
A concise revision guide to permanent dipole-permanent dipole interactions: which molecules have a permanent dipole, how polar molecules attract each other, why the attraction adds to London forces, and how to decide the forces in an unfamiliar molecule, for Edexcel International A Level Chemistry.
- 3.15-a
- 3.15-b
- 7.1ii
What these spec points say
- 3.15-a be able to distinguish between polar bonds and polar molecules
- 3.15-b be able to predict whether or not a given molecule is likely to be polar
- 7.1ii understand the nature of the following intermolecular forces: permanent dipole-permanent dipole interactions
Permanent Dipole-Dipole Forces
Choose a step to see why the H–Cl bond is polar, how neighbouring HCl molecules attract, how they line up in the liquid, and why a molecule with polar bonds can still have no permanent dipole.
© Dr. Mohammed Al-Fatah – onlinelearningsystem.net
GCSE Recap: Shared Pairs and Simple Molecules
Four quick GCSE questions on covalent bonds and molecules before you meet polar molecules.
Polar Molecules and Permanent Dipoles
A permanent dipole exists in a molecule that contains polar bonds whose dipoles do not cancel, so one end of the molecule is always δ+ and the other always δ-. Permanent dipole-permanent dipole interactions are the attractions between the δ+ end of one such molecule and the δ- end of a neighbour.
Two conditions must both hold. First, the molecule must contain a bond between atoms of different electronegativity, such as H-Cl, C-Cl, C-F or C=O. Second, the shape of the molecule must leave an overall dipole.
CO₂ has two very polar C=O bonds, but the molecule is linear and the two bond dipoles cancel. So CO₂ has no permanent dipole and only London forces act between its molecules.
The same applies to CCl₄ and BF₃.
| Molecule | Polar bonds? | Shape | Overall dipole? | Strongest intermolecular force |
|---|---|---|---|---|
| HCl | yes | linear, two atoms | yes | permanent dipole-permanent dipole interactions |
| CH₃Cl | yes | tetrahedral, one Cl | yes | permanent dipole-permanent dipole interactions |
| CO₂ | yes | linear, symmetric | no | London forces |
| CCl₄ | yes | tetrahedral, symmetric | no | London forces |
| SO₂ | yes | bent | yes | permanent dipole-permanent dipole interactions |
| N₂ | no | linear | no | London forces |
Key idea: Polar bonds are necessary but not sufficient. Ask whether the bond dipoles cancel in three dimensions before deciding that a molecule is polar.
Quick Check: Do the Bond Dipoles Cancel?
Four molecules, four shapes, one overall dipole. Draw them before you choose.
How the Molecules Line Up
In a liquid such as hydrogen chloride the molecules tumble constantly. On average, though, they spend more time with their δ+ hydrogen close to the δ- chlorine of a neighbour than in any other arrangement.
This partial alignment produces a net attraction that is stronger than London forces alone for molecules of similar size, so more energy is needed to pull the molecules apart.
Polar molecules attract each other by the alignment of their permanent dipoles; this attraction acts in addition to the London forces that every molecule has.
The alignment is never perfect, because thermal motion keeps disturbing it, which is why permanent dipole-permanent dipole interactions are still weak compared with covalent bonds and why polar liquids such as HCl still boil far below room temperature.
Remember: Permanent dipole-permanent dipole interactions act in addition to London forces, never instead of them.
Quick Check: Lined Up or Tumbling?
Judge one statement about what polar molecules are actually doing in a liquid.
Comparing Molecules With Similar Numbers of Electrons
The fair way to show that a permanent dipole raises the boiling temperature is to compare two molecules with the same number of electrons, so that their London forces are similar.
Hydrogen chloride and fluorine both have 18 electrons. F₂ is non-polar and boils at -188 °C; HCl is polar and boils at -85 °C.
Key idea: The extra 100 °C comes from the permanent dipole-permanent dipole interactions.
| Pair | Electrons | Polar? | boiling temperature / °C |
|---|---|---|---|
| F₂ | 18 | no | -188 |
| HCl | 18 | yes | -85 |
| Butane, C₄H₁₀ | 34 | no | -1 |
| Propanone, CH₃COCH₃ | 32 | yes (C=O) | 56 |
| Ethane, C₂H₆ | 18 | no | -89 |
| Fluoromethane, CH₃F | 18 | yes (C-F) | -78 |
Propanone and butane make the same point with organic molecules: almost the same electron count, but the C=O dipole in propanone lifts its boiling temperature by nearly 60 °C.
Exam sentence: Propanone has a higher boiling temperature than butane because propanone molecules have a permanent dipole, so permanent dipole-permanent dipole interactions act between them in addition to London forces, and more energy is needed to separate the molecules.
Quick Check: A Fair Comparison of Two Small Molecules
Build the comparison between two molecules of almost equal electron count with a 124 °C gap in boiling temperature.
Quick Check: Explain a Boiling Temperature
Write a short explanation for two isomers that differ only in where their chlorine atoms sit.
Deciding the Forces in an Unfamiliar Molecule
Exam questions often give a molecule you have never met and ask which intermolecular forces act between its molecules. Use three questions in order.
- Does the molecule contain an H atom bonded directly to N, O or F? If yes, hydrogen bonding is present (see the Hydrogen Bonding page), and so are the two forces below.
- Does the molecule have an overall permanent dipole? Check for polar bonds and then check that the shape does not cancel them. If yes, permanent dipole-permanent dipole interactions are present, plus London forces.
- Every molecule has London forces. If the answers to questions 1 and 2 are no, these are the only forces.
For example, chloromethane, CH₃Cl, has no H bonded to N, O or F. It has a polar C-Cl bond in a tetrahedral molecule with only one chlorine, so it has an overall dipole: permanent dipole-permanent dipole interactions plus London forces.
Methanal, HCHO, has a C=O dipole: the same answer.
Tetrachloromethane, CCl₄, has four polar bonds that cancel: London forces only.
Exam focus: Name every force that is present and say which is the strongest. A list with one force missing loses the mark even if the strongest is right.
Ask whether there is an H on N, O or F, then whether there is an overall dipole, and remember that every molecule also has London forces.
Quick Check: Name Every Force
Run the three questions in order over five molecules you have not met on this page.
Permanent Dipoles and Solubility
Polar molecules tend to dissolve in polar solvents and non-polar molecules in non-polar solvents. A solute dissolves when the new attractions it makes with solvent molecules are similar in strength to the attractions it loses.
Propanone mixes with water in all proportions because its C=O oxygen can accept a hydrogen bond from water.
Hexane does not mix with water because the only forces hexane can offer are London forces. The Choosing Solvents page develops this idea.
Key idea: Like dissolves like: match the strongest intermolecular force in the solute to the strongest force in the solvent.
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 HCl boils higher than F₂
Similar electrons, so similar London forces; HCl is polar, so permanent dipole-permanent dipole interactions act as well; more energy is needed to separate HCl molecules.
Explain why CO₂ is non-polar although C=O bonds are polar
The molecule is linear and symmetrical, so the two bond dipoles cancel and there is no overall dipole.
State the intermolecular forces in CH₃Br
Permanent dipole-permanent dipole interactions and London forces; no hydrogen bonding because no H is bonded to N, O or F.
Do not say
“Polar molecules have no London forces”; “any molecule with a polar bond is polar”.
FAQs
Use these quick answers to check how Edexcel International expects permanent dipoles to be identified and explained.
What is a permanent dipole?
It is a separation of charge that is always present in a molecule, because the molecule contains polar bonds whose dipoles do not cancel. One end of the molecule is permanently δ+ and the other permanently δ-.
Does a polar bond always make a polar molecule?
No. CO₂ and CCl₄ both contain very polar bonds, but the molecules are symmetrical, so the bond dipoles cancel and there is no overall dipole. Always check the shape before deciding.
How do I decide which forces act between molecules of an unfamiliar substance?
Ask three questions in order: is a hydrogen atom bonded directly to N, O or F (hydrogen bonding); does the molecule have an overall permanent dipole (permanent dipole-permanent dipole interactions); and finally, every molecule has London forces.
Which is stronger, a permanent dipole attraction or London forces?
For molecules of a similar size, the permanent dipole-permanent dipole attraction is stronger. For a much larger molecule the London forces can be bigger in total, which is why HI boils higher than HCl even though HCl is more polar.
How should I show this in an answer?
Compare molecules with a similar number of electrons so the London forces are similar, then attribute the difference to the permanent dipole and finish with the energy statement.
Copyright notice: This OLS revision content, including the explanations, layout, diagrams, tables and embedded learning structure, is authored for Online Learning System by Dr. Mohammed Al-Fatah. It may not be copied, reproduced, redistributed or adapted without written permission.
Keep this note — free
Save your progress across every Edexcel IAL topic. A free account remembers which topics you have covered, saves your question scores, and syncs across your phone and laptop.
- Track every topic you have finished
- Keep your practice-question scores
- No payment, no card, free forever
Already registered? Log in