Intermolecular Forces and States of Matter Study Pack
Kibin's free study pack on Intermolecular Forces and States of Matter includes a 6-section study guide, 26 quiz questions, 30 flashcards, and 5 open-ended Explain review questions. Sign up free to track your progress toward mastery, plus upload your own notes and recordings to create personalized study packs organized by course.
Last updated May 27, 2026
Intermolecular Forces and States of Matter Study Guide
Unpack the full hierarchy of intermolecular forces — from London dispersion and dipole–dipole interactions to hydrogen bonding and ion–dipole forces — and see how each shapes boiling points, viscosity, and states of matter.
Key Takeaways
- •Intermolecular forces (IMFs) are electrostatic attractions between molecules that determine physical properties such as boiling point, melting point, viscosity, and surface tension.
- •Dispersion forces (London dispersion) arise from temporary, instantaneous dipoles and act between all molecules; their strength increases with molecular size and polarizability.
- •Dipole–dipole forces act between polar molecules whose permanent unequal charge distributions create attractive interactions between partially positive and partially negative ends.
- •Hydrogen bonding is a uniquely strong dipole–dipole interaction that occurs when hydrogen is covalently bonded to nitrogen, oxygen, or fluorine and attracted to a lone pair on one of those same electronegative atoms in a neighboring molecule.
- •The physical state of a substance — solid, liquid, or gas — reflects the balance between the kinetic energy of particles and the strength of the intermolecular forces holding them together.
- •Ion–dipole forces, the strongest of all IMFs, occur between an ion and a polar molecule and are critical to the dissolution of ionic compounds in polar solvents like water.
- •Relative IMF strength follows the general order: ion–dipole > hydrogen bonding > dipole–dipole > London dispersion, though London dispersion forces can dominate in large, highly polarizable nonpolar molecules.
What Intermolecular Forces Are and Why They Matter
Intermolecular forces are electrostatic attractions and repulsions that act between separate molecules (or between molecules and ions), distinct from the intramolecular covalent or ionic bonds that hold atoms together within a single molecule. These forces govern whether a substance is a gas, liquid, or solid under a given set of conditions, and they directly control measurable physical properties.
Intermolecular vs. Intramolecular Interactions
- •Intramolecular bonds — covalent and ionic bonds — hold atoms together within a molecule and are much stronger than IMFs.
- •Intermolecular forces act between molecules; breaking them does not change molecular identity, only physical state.
- •Boiling a liquid breaks IMFs, not covalent bonds — this is why water's covalent O–H bonds remain intact when water vaporizes.
Physical Properties Controlled by IMF Strength
- •Boiling point and melting point both increase as IMF strength increases because more thermal energy is required to overcome stronger attractions.
- •Viscosity (resistance to flow) and surface tension both rise with stronger IMFs because molecules resist being separated or pulled away from the bulk liquid.
- •Vapor pressure decreases as IMF strength increases, since fewer molecules have enough energy to escape the liquid phase into the gas phase.
London Dispersion Forces: IMFs Between All Molecules
London dispersion forces (also called dispersion forces or van der Waals forces in the narrow sense) are the only type of intermolecular force present in nonpolar molecules, though they operate between every molecule regardless of polarity. They arise from the quantum-mechanical reality that electron density is never perfectly static.
Origin of Instantaneous and Induced Dipoles
- •At any given instant, the electrons in a molecule may be unevenly distributed, creating a brief instantaneous dipole — a momentary separation of partial positive and partial negative charge.
- •This instantaneous dipole induces a complementary dipole in a neighboring molecule, and the two temporary dipoles attract each other.
- •These fluctuations occur constantly, producing a net attractive force even though both molecules are nonpolar on average.
Factors That Increase London Dispersion Strength
- •Polarizability — the ease with which an electron cloud can be distorted — increases with the number of electrons and the size of the electron cloud.
- •Larger, heavier molecules (higher molar mass) have more electrons and greater polarizability, producing stronger London dispersion forces and higher boiling points.
- •Molecular shape also matters: elongated molecules like n-pentane have greater surface contact area and stronger dispersion forces than compact, spherical molecules like neopentane (same molecular formula, lower boiling point).
Dipole–Dipole Forces and Hydrogen Bonding in Polar Molecules
Polar molecules experience additional IMFs beyond London dispersion because their permanent, unequal charge distributions create persistent partial charges at specific ends of the molecule. The most important special case of dipole–dipole interaction is hydrogen bonding, which is strong enough to produce dramatic, observable effects on physical properties.
Dipole–Dipole Attractions in Polar Molecules
- •A polar molecule has a permanent dipole moment resulting from differences in electronegativity between bonded atoms, leaving one end of the molecule with a partial positive charge (δ+) and the other with a partial negative charge (δ−).
- •In the liquid or solid phase, polar molecules orient so that δ+ regions of one molecule face δ− regions of neighbors, producing an electrostatic attraction.
- •Dipole–dipole forces are stronger than London dispersion forces for molecules of comparable size, which is why polar molecules generally have higher boiling points than nonpolar molecules of similar molar mass.
Hydrogen Bonding: Conditions and Strength
- •Hydrogen bonding requires hydrogen to be covalently bonded to one of three highly electronegative atoms — nitrogen (N), oxygen (O), or fluorine (F) — making the hydrogen strongly δ+.
- •The δ+ hydrogen is then attracted to a lone pair of electrons on an N, O, or F atom of a neighboring molecule.
- •This interaction is unusually strong among IMFs (roughly 5–25 kJ/mol) because N, O, and F are so electronegative and small that the partial charges are intense and the interacting atoms get very close.
- •Hydrogen bonding explains water's anomalously high boiling point (100 °C) compared to other small hydrides like H₂S (−60 °C), as well as ice's lower density than liquid water — the open hexagonal hydrogen-bonded network in ice is less dense than the more disordered liquid.
- •DNA base pairing and protein secondary structure (α-helices and β-sheets) are stabilized by hydrogen bonding, making it biologically indispensable.
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What is the correct order of intermolecular force strength from strongest to weakest?
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Concept 1 of 5
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Intermolecular Forces vs. Intramolecular Bonds
Explain the difference between intermolecular forces and intramolecular bonds in your own words. Why is this distinction important for understanding physical properties like boiling point?
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