Intermolecular forces are attractions between separate molecules; they are different from the covalent or ionic bonds that hold atoms together within a molecule or solid. These attractions help explain boiling points and other properties, but their effects depend on the molecules involved—not just on a simple ranking of force names.
Intermolecular forces versus intramolecular bonds
In a water molecule, covalent bonds join its hydrogen and oxygen atoms. Intermolecular forces act between one water molecule and another. The distinction matters during a phase change: when liquid water boils, molecules separate into the gas phase, but the water molecules ordinarily remain intact. Vaporization overcomes attractions between molecules; it does not ordinarily break the molecules’ covalent bonds.
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OpenStax illustrates the difference with hydrogen chloride. In its textbook example, vaporizing one mole of liquid HCl takes about 17 kJ, while breaking the H–Cl covalent bonds in one mole takes about 430 kJ. These figures describe that educational example, not universal energies for every substance or condition. OpenStax, Chemistry: Atoms First 2e, section 10.1
How charge distribution creates attractions
Electrons are constantly moving, and their distribution can be uneven. The resulting charge patterns explain several common intermolecular interactions. Introductory chemistry often focuses on London dispersion, dipole–dipole attraction, and hydrogen bonding. That classroom grouping is useful, but it is not an exhaustive taxonomy: ion–dipole and other multipole interactions also occur.
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London dispersion: present in every molecule
A momentary uneven electron distribution can create a temporary dipole, which can induce a dipole in a nearby particle. The resulting attraction is called London dispersion. It occurs in all molecules, including nonpolar ones; a molecule does not need a permanent dipole to experience dispersion.
Dispersion often becomes more influential as particles grow larger and more polarizable, because their electron clouds are easier to distort. OpenStax’s textbook boiling-point values for the halogens illustrate the trend: F2, 85 K; Cl2, 238 K; Br2, 332 K; and I2, 457 K. This series illustrates how increasing size and polarizability can accompany stronger dispersion effects; it should not be taken to mean dispersion is the only factor in every boiling-point comparison. OpenStax, Chemistry: Atoms First 2e, section 10.1
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Dipole–dipole attraction: permanent polarity matters
Polar molecules have a permanent separation of partial charge. When the partially positive end of one molecule faces the partially negative end of another, they attract. Their orientation matters: other arrangements of dipoles can be less attractive or even repulsive.
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Repair common Windows errors and clear accumulated junk for a smoother, more stable PC - no reinstall needed.Free scan · no reinstallHCl is polar, so its molecules experience dipole–dipole attraction as well as dispersion. F2 is nonpolar but still experiences dispersion. OpenStax gives normal boiling points of 188 K for HCl and 85 K for F2; the similar-mass comparison illustrates how permanent polarity can affect boiling point. These are textbook example values, not a general prediction that polarity alone determines boiling point. OpenStax, Chemistry: Atoms First 2e, section 10.1
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Hydrogen bonding: a directional, strong dipole interaction
In the introductory framework, hydrogen bonding commonly involves hydrogen covalently attached to N, O, or F interacting with an electron-rich site on another molecule. The interaction is particularly strong and geometrically meaningful compared with many other dipole interactions, but it is not a covalent bond joining the two molecules.
Water, hydrogen fluoride, and ammonia are familiar examples. Hydrogen bonding helps account for water being liquid at ordinary temperatures and also matters in biological structures: it contributes to enzyme shape and DNA base pairing. OpenStax, Organic Chemistry, section 2.12 OpenStax’s Chemistry: Atoms First 2e states: “Hydrogen bonds are intermolecular attractive forces, not intramolecular attractive forces (covalent bonds).” OpenStax, section 10.1
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Which intermolecular forces are present?
Use this practical sequence when identifying interactions between molecules:
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- Check for permanent polarity. If a molecule is polar, its molecules can have dipole–dipole interactions. Consider their orientations, since not every alignment is attractive.
- Check for hydrogen-bonding partners. In the common introductory definition, look for hydrogen attached to N, O, or F and an electron-rich site on a neighboring molecule. Assess whether the molecules can donate and accept hydrogen bonds under the definition being used.
- Include dispersion in every case. It is present whether the molecules are polar or nonpolar.
- Consider the broader setting. If ions are present, ion–dipole interactions may matter; more advanced treatments also account for multipole interactions.
MIT OpenCourseWare’s examples include HCl for dipole–dipole interactions; argon, iodine, methane, and helium for induced-dipole interactions; and HF, water, and ammonia for hydrogen bonding. MIT OpenCourseWare, “12. Intermolecular Forces,” Fall 2010
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How intermolecular forces affect boiling points
Boiling requires molecules to separate into the gas phase. Stronger attractions can mean more energy is needed to separate them, influencing boiling points. But a boiling point is not a direct measurement of one named force: molecular size, polarizability, shape and contact area, as well as the other interactions present, all affect comparisons.
A useful comparison starts with polarity and possible hydrogen bonding, then considers size and polarizability, shape and contact area, and the property or conditions at issue. Keep two ideas separate: the strongest named interaction a substance has, and the largest contribution to its total intermolecular attraction. In more complex comparisons, those need not be the same.
Is “van der Waals force” the same as intermolecular force?
Usage varies. OpenStax uses “van der Waals” broadly for attractions between neutral atoms and molecules, then discusses component types. The IUPAC educational overview describes dispersion and multipole interactions among neutral molecules, treats hydrogen bonding and ion–dipole interactions as distinct categories, and notes that “van der Waals” is often used for dispersion plus dipole and other multipole interactions. For a broad, clear label, “intermolecular forces” or “intermolecular interactions” avoids relying on one convention. IUPAC, “Types of intermolecular forces”
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Further study
For an introductory explanation, see OpenStax Chemistry: Atoms First 2e, section 10.1. MIT OpenCourseWare’s lesson also lists Jacob Israelachvili’s Intermolecular and Surface Forces (Academic Press, 1992) as supplemental reading for a more advanced treatment. MIT OpenCourseWare, Fall 2010 lesson
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