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Intermolecular forces are attractions between separate molecules; they are not the covalent bonds that hold atoms together within a molecule. These attractions help explain properties such as boiling point, but their effects depend on the molecules involved and the conditions—not on a single universal ranking of force types.
Intermolecular forces versus intramolecular bonds
Intramolecular bonds join atoms within a molecule. Covalent bonds, for example, involve atoms sharing electrons. Intermolecular forces act between distinct molecules. The phrase “intermolecular bonds” is common in teaching, but “forces” or “interactions” is often more precise because these attractions do not ordinarily join molecules into one covalently bonded unit.
A phase change such as vaporization usually separates molecules by overcoming their attractions to one another; it does not ordinarily break their covalent bonds. OpenStax’s Chemistry: Atoms First 2e illustrates the difference with HCl: its example uses about 17 kJ per mole to vaporize liquid HCl, compared with about 430 kJ per mole to break the H–Cl covalent bonds in one mole of HCl. These are textbook example values, not universal energies for every sample or condition. OpenStax, “10.1 Intermolecular Forces”.
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Electrons are distributed around nuclei, and that distribution determines whether a molecule has a permanent dipole or can form temporary ones. Molecular polarity, electron-cloud flexibility (polarizability), and the way molecules approach one another all matter.
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London dispersion: present in every molecule
Electrons move, so a molecule’s electron distribution can fluctuate and create a temporary dipole. That temporary dipole can induce a dipole in a nearby particle, producing an attractive interaction called London dispersion. Dispersion occurs in nonpolar molecules as well as polar ones; a molecule does not need a permanent dipole to experience it. MIT OpenCourseWare uses Ar, I2, CH4, and He as examples of particles with induced-dipole interactions. MIT OpenCourseWare, “12. Intermolecular Forces”.
Dipole-dipole attraction: permanent polarity matters
Polar molecules have regions of partial positive and partial negative charge. When molecules orient so that opposite partial charges face one another, they attract. Their orientation can also put like partial charges near one another, creating repulsion; “dipole-dipole” does not mean every orientation is attractive. HCl is a familiar example of a polar molecule with dipole-dipole attraction, alongside dispersion.
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Hydrogen bonding: a directional, relatively strong dipole interaction
In the introductory chemistry framework, hydrogen bonding commonly involves hydrogen covalently attached to N, O, or F interacting with an electron-rich site on another molecule. Water, HF, and ammonia are standard examples. Hydrogen bonding is an intermolecular attraction, not a separate covalent bond between molecules. OpenStax states: “Hydrogen bonds are intermolecular attractive forces, not intramolecular attractive forces (covalent bonds).” Its organic chemistry text connects hydrogen bonding with water’s liquid state at ordinary temperatures, enzyme shape, and DNA base pairing. OpenStax, “2.12 Noncovalent Interactions between Molecules”.
Are there only three kinds of intermolecular force?
Dispersion, dipole-dipole interactions, and hydrogen bonding are a useful introductory set, not an exhaustive taxonomy. Broader treatments also discuss ion-dipole interactions and multipole interactions. Terminology varies: OpenStax uses “van der Waals forces” broadly for attractions between neutral atoms and molecules, then describes component types. IUPAC’s educational overview treats hydrogen bonding and ion-dipole interactions as distinct categories and notes that “van der Waals” is often used for dispersion and dipole or other multipole interactions. For general use, “intermolecular forces” or “intermolecular interactions” is the clearest umbrella term. IUPAC, “Types of intermolecular forces”.
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How do intermolecular forces affect boiling point?
Boiling requires molecules to separate into the gas phase. Stronger overall attractions can therefore contribute to a higher boiling point, but a boiling-point comparison is not a simple measurement of one named force. Molecular size and polarizability, shape and contact area, the other interactions present, and the conditions being compared can all affect the outcome.
HCl and F2: a polarity comparison
OpenStax gives normal boiling points of 188 K for HCl and 85 K for F2. The molecules have similar masses, making this a useful textbook illustration of how HCl’s permanent polarity and resulting dipole-dipole attraction can contribute to a higher boiling point. Both substances also experience dispersion; the comparison does not mean dispersion is absent from F2. OpenStax, “10.1 Intermolecular Forces”.
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The halogens: size and polarizability
OpenStax’s textbook values show increasing boiling points down this series of nonpolar diatomic molecules:
| Substance | Textbook normal boiling point |
|---|---|
| F2 | 85 K |
| Cl2 | 238 K |
| Br2 | 332 K |
| I2 | 457 K |
The trend illustrates how greater molecular size and polarizability can strengthen dispersion. It should not be read as proof that dispersion alone determines every boiling point. These are educational example values from OpenStax, “10.1 Intermolecular Forces”.
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How to identify the interactions between two substances
- Check polarity. Determine whether each molecule has a permanent dipole. If it does, dipole-dipole interactions are possible; all molecules also experience dispersion.
- Check for hydrogen bonding. In the introductory convention, look for hydrogen bonded to N, O, or F and an electron-rich site that can interact with it on another molecule. Definitions and advanced treatments can be more detailed.
- Compare size and polarizability. Larger, more easily distorted electron clouds can increase dispersion, so do not ignore molecule size when comparing substances.
- Consider shape and contact. Molecular shape affects how closely particles can approach and how much surface can interact.
- Specify what you are comparing. A question about the strongest named interaction present is different from a question about the largest total contribution to a property such as boiling point. Include the conditions and property of interest when making the comparison.
Which interaction is strongest?
There is no reliable universal ranking that predicts every pair of substances. Hydrogen bonding is often described as a particularly strong, geometrically meaningful dipole interaction in introductory chemistry, but the total attraction in a real comparison also reflects dispersion, molecular size and shape, orientation, and the environment. A small molecule with hydrogen bonding does not automatically have a higher boiling point than every larger molecule without it. Identify which interactions are possible, then compare the actual molecules and the property in question.
For a more advanced treatment, MIT OpenCourseWare lists Jacob Israelachvili’s Intermolecular and Surface Forces (Academic Press, 1992) as supplemental reading for its Fall 2010 lesson; it is not required for an introductory understanding.
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