Dipole-Dipole Forces and Ion-Dipole Interactions
Dipole–dipole forces are attractions between the positive end of one polar molecule and the negative end of another, while ion–dipole forces are attractions between an ion and a polar molecule, and they are much stronger. In brief: acetone (polar, 58 g/mol) boils at 56 °C, while butane (nonpolar, 58 g/mol) boils at −0.5 °C, and a sodium ion in water is held by about six water molecules with a hydration enthalpy of roughly −405 kJ/mol.
These two kinds of electrostatic attraction explain why polar substances have higher boiling points than nonpolar ones of similar mass and why ionic compounds dissolve in water. This article describes how the forces arise, compares their strengths with other interactions, and works through examples that show how to predict which substance has stronger attractions.
What are dipole–dipole forces?
A polar molecule has a permanent dipole, with a partial negative charge at one end and a partial positive charge at the other. In a liquid or solid, the molecules orient themselves so that the positive end of one is close to the negative end of the next, which lowers the energy of the system. Typical dipole–dipole forces are of the order of 5 to 25 kJ/mol. They are much weaker than covalent bonds (150–1000 kJ/mol).
The strength depends on the size of the dipole moments and on the distance. For molecules of similar size and shape, those with larger dipole moments attract each other more strongly. In a gas the molecules rotate freely, so the average attraction is much weaker than in a liquid and falls with distance as 1/r6.
What are ion–dipole interactions?
An ion has a full charge, which attracts the oppositely charged end of a polar molecule. A cation such as Na+ attracts the oxygen end of a water molecule, and an anion such as Cl− attracts the hydrogen end. In the dissolution of an ionic compound in water, these interactions replace the strong ion–ion attractions in the crystal.
The strength depends on the charge of the ion, on its size and on the dipole moment of the solvent molecule. Ions with higher charge and smaller radius attract the surrounding molecules more strongly, so Mg2+ is hydrated much more strongly than Na+.
How strong is hydration?
| Ion | Charge | Ionic radius (pm) | Hydration enthalpy (kJ/mol) |
|---|---|---|---|
| K+ | +1 | 138 | about −320 |
| Na+ | +1 | 102 | about −405 |
| Li+ | +1 | 76 | about −520 |
| Ca2+ | +2 | 100 | about −1600 |
| Mg2+ | +2 | 72 | about −1920 |
The values are approximate, and tables from different sources differ by a few percent. They show two trends: hydration becomes stronger for smaller ions, and much stronger for doubly charged ions. The large energy released helps to pay for the energy needed to break up the crystal lattice, so many salts dissolve despite their strong ionic bonds.
How do the forces compare in strength?
| Interaction | Between | Typical energy (kJ/mol) |
|---|---|---|
| Ion–ion | Oppositely charged ions | 400–4000 |
| Ion–dipole | Ion and polar molecule | 40–600 |
| Hydrogen bond | H–F, H–O, H–N and a lone pair | 10–40 |
| Dipole–dipole | Polar molecules | 5–25 |
| London dispersion | All particles | 0.05–40 |
These ranges overlap and vary in the literature, so they should be taken as orders of magnitude.
Worked examples
Example 1: acetone and butane
Acetone, (CH3)2C=O, has a molar mass of 58.08 g/mol and a dipole moment of 2.9 D. Butane, C4H10, has a molar mass of 58.12 g/mol and is nonpolar. With almost identical masses and thus similar dispersion forces, the difference in boiling point (56.1 °C against −0.5 °C) is mostly due to the dipole–dipole forces in acetone.
Example 2: choosing the stronger attraction
Which has the higher boiling point, CH3Cl (−24.2 °C) or CH3CH3 (−88.6 °C)? Methyl chloride is polar (1.87 D) and has more electrons, so it has stronger attractions and a higher boiling point. Ethane has a similar mass (30 g/mol compared with 50.5 g/mol) but no dipole.
Example 3: identifying the forces in a solution
What forces act when sodium chloride dissolves in water? In the solution, ion–dipole forces hold each Na+ to the oxygen atoms of water, and each Cl− to hydrogen atoms; dipole–dipole forces (with hydrogen bonds) hold water molecules to each other, and dispersion forces act between all of them.
Example 4: dissolving a salt in a nonpolar solvent
NaCl does not dissolve in hexane, because hexane molecules are nonpolar and can offer only weak dispersion forces, which are not strong enough to replace the strong ion–ion attractions of the crystal.
Where do these forces matter?
- Electrolyte solutions: ion–dipole attractions make seawater, blood plasma and cell fluids conduct electricity.
- Ion channels: proteins in cell membranes use polar groups to replace the water shell of an ion, so they can let K+ pass but not Na+.
- Boiling and melting points: polar organic compounds such as ketones and nitriles boil higher than hydrocarbons of similar size.
- Chromatography: polar compounds stick more strongly to polar stationary phases such as silica and are separated from nonpolar ones.
How were the forces understood?
Peter Debye’s work on dipole moments in the 1910s and the work of Willem Keesom in 1912 on the average attraction between rotating permanent dipoles formed the basis of the theory of dipole–dipole forces. The Keesom interaction is therefore a name for the dipole–dipole force in physical chemistry. The study of hydration of ions started with Svante Arrhenius’s theory of dissociation in 1884, and later Max Born derived in 1920 an equation for the energy of an ion in a polar solvent.
Which mistakes are common?
- Assuming that dipole–dipole forces always dominate: dispersion forces can be larger in big molecules.
- Mixing up ion–dipole with ion–ion: the first involves a polar molecule, the second two ions.
- Forgetting that hydrogen bonds are a special strong form of dipole–dipole attraction: they are treated in their own right.
- Comparing substances of very different mass: compare similar masses to isolate the effect of polarity.
- Treating the energy ranges as exact: they overlap.