# Intermolecular force **Intermolecular forces** are the forces of attraction and repulsion that act between [[Molecule|molecules]], or between a molecule and a neighbouring [[Ion|ion]], as distinct from the [[Chemical_bond|chemical bonds]] that hold the atoms of one molecule together. They are weak by the standard of a [[Covalent_bond|covalent bond]] — breaking the bond in H₂ costs 436 kJ/mol, while the strongest common intermolecular attraction, the [[Hydrogen_bond|hydrogen bond]], is worth a few tens of kJ/mol — yet they decide whether a substance is a gas, a liquid or a solid at a given [[Temperature|temperature]], how high its [[Boiling_point|boiling point]] is, how readily it flows and how it wets a surface.[^cboc-h2][^iupac-hb] The family divides into hydrogen bonding, [[Dipole|dipole]]–dipole and ion–dipole attractions, and the three [[Van_der_Waals_force|van der Waals forces]] named for Keesom, Debye and London, of which the London dispersion force acts between all molecules whatever their shape or charge. In the microsim below the reader picks one of four groups of the [[Periodic_table|periodic table]] (14, 15, 16 or 17) and reads the boiling points of its four simple hydrides against period: CH₄ to SnH₄, NH₃ to SbH₃, [[Water|H₂O]] to H₂Te, HF to HI. Down each group the dispersion trend climbs with molar mass; in three of the four groups the period-2 member sits far above that trend, and the readout gives the anomaly in kelvin together with the strongest force acting between each pair of molecules (dispersion, dipole–dipole or hydrogen bond). The question the sim answers is the pair's own: why water is a liquid at room temperature while hydrogen sulfide, its heavier cousin, is a gas. On the [[Chemistry]] flagship this article is the child of Part IV — Bonding, section *Intermolecular forces* (row K21), between the bonding rows and the Part V energy rows, where the same attractions reappear as the enthalpy of vaporization measured by [[Calorimetry|calorimetry]]. ## Hydrogen bonding A hydrogen bond forms when a hydrogen atom covalently bound to a strongly electronegative atom — in practice [[Nitrogen|nitrogen]], [[Oxygen|oxygen]] or [[Fluorine|fluorine]] — is attracted to a lone pair on a second electronegative atom nearby. IUPAC's 2011 recommendation defines it more generally as an attractive interaction between a hydrogen atom from a molecule or fragment X–H, in which X is more electronegative than H, and an atom or group of atoms in the same or a different molecule, where there is evidence of bond formation.[^iupac-hb] The [[Electronegativity|electronegativity]] difference leaves the hydrogen with a large partial positive charge and almost no electron density of its own, so it can approach the acceptor's lone pair closely, and the interaction is directional in a way that ordinary dipole–dipole attraction is not.[^iupac-hb] The microsim's four ladders put a number on the effect. Each ladder plots the normal boiling point of the four hydrides in one group against period. In the dispersion-only groups the points rise steadily with molar mass; the sim draws a straight line through the period-3, -4 and -5 members and extrapolates it back to period 2 (an ILLUSTRATIVE display fit, not a law), then reports the gap between that line and the measured value. For group 16 the line lands near 185 K, and water boils at 373 K: an anomaly of roughly 190 K, the largest of the four. For group 17 hydrogen fluoride sits about 130 K above its line, for group 15 [[Ammonia|ammonia]] about 90 K, and for group 14 methane sits close to its line, because [[Carbon|carbon]] is not electronegative enough to donate a hydrogen bond and has no lone pair to accept one.[^af2e-10-1][^nist-webbook] The readout names the force at work: dispersion for CH₄ and the heavier members of every group, dipole–dipole for the polar hydrides below period 2, and hydrogen bonding for H₂O, HF and NH₃. Water's anomaly is the largest because each molecule can donate two hydrogen bonds and accept two, so liquid water is a three-dimensional network rather than a set of chains, as in HF, or a looser web, as in ammonia.[^af2e-10-1] The same network explains why [[Ice|ice]] floats: the fully hydrogen-bonded solid is an open hexagonal lattice and is less dense than the liquid it melts into.[^af2e-10-1] In [[Biochemistry|biochemistry]] hydrogen bonds hold the two strands of DNA together through the base pairs proposed by Watson and Crick in 1953, and they stabilise the helices and sheets of protein structure.[^watson-crick] ## Salt bridge A salt bridge is a pair of oppositely charged groups held together at close range, usually an ionic attraction combined with a hydrogen bond. In proteins the term is used for the contact between a positively charged side chain (lysine or arginine) and a negatively charged one (aspartate or glutamate); Barlow and Thornton's 1983 survey of ion pairs in protein crystal structures adopted a distance criterion of 4 Å or less between the charged atoms, and that cutoff is still the usual working definition.[^barlow1983] Salt bridges contribute to the folding and thermal stability of proteins and to the recognition between an [[Enzyme|enzyme]] and its substrate. The electrostatic part follows [[Coulomb's_law|Coulomb's law]] and is screened by the surrounding water and dissolved [[Electrolyte|electrolyte]], so a salt bridge buried in the protein interior is worth more than one exposed on the surface.[^kumar2002] Designed receptors in [[Supramolecular_chemistry|supramolecular chemistry]] bind carboxylate or phosphate guests through the same pairing. ## Dipole–dipole and similar interactions A polar molecule carries a permanent [[Electric_dipole_moment|electric dipole moment]] because its shared electrons sit closer to the more electronegative atoms; the [[Chemical_polarity|polarity]] of the whole molecule then depends on its [[Molecular_geometry|geometry]], which is why linear CO₂ is non-polar while bent H₂O is not. Two such dipoles attract when they align head to tail and repel when they align head to head. In a liquid the molecules tumble, so the net effect is a thermal average over orientations that favours the attractive arrangements; the averaged energy falls off as the inverse sixth power of the separation and weakens as the temperature rises, which is the Keesom form discussed below. For the sim's hydrides the dipole–dipole contribution is what lifts PH₃, H₂S and HCl slightly above the non-polar group 14 hydride of the same period, while the far larger lift in period 2 belongs to hydrogen bonding.[^af2e-10-1] ### Ion–dipole and ion–induced dipole forces An ion sitting next to a polar molecule attracts the end of the dipole that carries the opposite partial charge. This ion–dipole force is stronger than dipole–dipole attraction because a full charge is involved, and it is the reason [[Sodium_chloride|sodium chloride]] dissolves in water: each Na⁺ is surrounded by the oxygen ends of water molecules and each Cl⁻ by their hydrogen ends, and the energy released by this hydration repays most of the [[Lattice_energy|lattice energy]] of the crystal, 787 kJ/mol in the Born–Haber accounting.[^af2e-10-1][^cboc-nacl] An ion can also pull the electron cloud of a non-polar molecule toward or away from itself, creating a temporary dipole in a molecule that had none; this ion–induced dipole force is weaker and shorter-ranged, and it is part of what holds a non-polar molecule such as O₂ near a metal ion or in an ionic [[Solution_(chemistry)|solution]]. ## Van der Waals forces The van der Waals forces are the attractions and repulsions between molecules that arise neither from a full charge nor from a hydrogen bond. They are named for Johannes Diderik van der Waals, whose 1873 Leiden thesis on the continuity of the gas and liquid states introduced an [[Van_der_Waals_equation|equation of state]] with one correction for intermolecular attraction and another for molecular volume, work recognised with the 1910 Nobel Prize in Physics.[^vdw-nobel] In the modern classification the attractive part is the sum of three contributions, each varying as the inverse sixth power of the distance between the molecules, set out in the next three sections; at very short range they are opposed by a steep repulsion when the electron clouds overlap, and the [[Lennard-Jones_potential|Lennard-Jones potential]] of 1924, with its r⁻¹² wall and r⁻⁶ tail, is the usual compact description of the pair.[^lj1924] ### Keesom force (permanent dipole – permanent dipole) The Keesom force is the thermally averaged attraction between two permanent dipoles, worked out by Willem Hendrik Keesom in 1921.[^keesom1921] For two molecules with dipole moments μ₁ and μ₂ a distance r apart, the orientation-averaged energy is `U = −(2/3)·μ₁²·μ₂²/((4πε₀)²·k_B·T·r⁶)`, where k_B is the [[Boltzmann_constant|Boltzmann constant]]: the energy is negative (attractive), falls as r⁻⁶, and shrinks as temperature rises because thermal motion randomises the orientations. It is the only one of the three van der Waals terms that depends on temperature, and it requires both partners to be polar. ### Debye force (permanent dipoles–induced dipoles) The Debye force, described by Peter Debye in 1920, arises when the field of a permanent dipole induces a dipole in a neighbouring molecule, polar or not.[^debye1920] The induced moment is proportional to the neighbour's polarizability α, and the energy, `U = −μ²·α/((4πε₀)²·r⁶)`, is again attractive and again r⁻⁶, but it does not average away with temperature because the induced dipole always follows the inducing one. It is usually the smallest of the three terms in a polar liquid. ### London dispersion force (fluctuating dipole–induced dipole interaction) The London dispersion force acts between all atoms and molecules, including the [[Noble_gas|noble gases]] and non-polar molecules such as CH₄, and is the reason those substances condense at all. Fritz London explained it in 1930 with [[Quantum_mechanics|quantum mechanics]]: the electrons of one molecule are never at rest, so at any instant it carries a fleeting dipole that polarises its neighbour, and the two instantaneous dipoles are correlated so that their attraction survives the average.[^london1930] London's result for two identical molecules is `U ≈ −(3/4)·I·α²/((4πε₀)²·r⁶)`, with I the [[Ionization_energy|ionization energy]] and α the polarizability.[^london1937] Because the polarizability grows with the number of electrons and the size of the electron cloud, the dispersion force grows down a group of the periodic table, and this is the trend the microsim's baseline follows: SnH₄ boils at 221 K against 161 K for SiH₄, and H₂Te at 271 K against 214 K for H₂S.[^nist-webbook] Shape matters too, since elongated molecules present more surface to their neighbours than compact ones. Summed over a large enough contact area dispersion can outweigh hydrogen bonding, and it is dispersion, acting across millions of fine hairs, that lets a gecko hold its weight on glass.[^autumn2002] ## Relative strength of forces The table sets the microsim's data beside the bond energies the forces are usually compared with. The covalent bond in H₂ costs 436 kJ/mol to break and the [[Ionic_bonding|ionic]] lattice of NaCl 787 kJ/mol per formula unit; hydrogen bonds are typically a few tens of kJ/mol, dipole–dipole attractions a few kJ/mol, and dispersion anything from a fraction of a kJ/mol for [[Helium|helium]] to tens of kJ/mol for large molecules.[^cboc-h2][^cboc-nacl][^cboc-ch6] The [[Enthalpy_of_vaporization|enthalpy of vaporization]] is the practical measure, because boiling breaks intermolecular contacts and leaves the covalent bonds alone: water needs 40.7 kJ/mol at its boiling point, methane 8.2 kJ/mol.[^nist-webbook] | Group | Period 2 | Period 3 | Period 4 | Period 5 | Strongest force in period 2 | |---|---|---|---|---|---| | 14 | CH₄ 112 K | SiH₄ 161 K | GeH₄ 185 K | SnH₄ 221 K | dispersion | | 15 | NH₃ 240 K | PH₃ 185 K | AsH₃ 211 K | SbH₃ 256 K | hydrogen bond | | 16 | H₂O 373 K | H₂S 214 K | H₂Se 232 K | H₂Te 271 K | hydrogen bond | | 17 | HF 293 K | HCl 188 K | HBr 207 K | HI 238 K | hydrogen bond | *Normal boiling points of the group 14–17 hydrides, rounded to the nearest kelvin: the microsim's data table.*[^nist-webbook][^af2e-10-1] Reading down any column the values climb with molar mass, which is dispersion; reading across the period-3 row from SiH₄ to HCl the values rise only modestly, which is the dipole–dipole increment. Reading across period 2 the sequence CH₄ < NH₃ < HF < H₂O is neither the order of molar mass nor the order of electronegativity — fluorine is more electronegative than oxygen, yet HF boils 80 K below water — but the order of how many hydrogen bonds each molecule can make: HF has three lone pairs and one hydrogen, NH₃ three hydrogens and one lone pair, and only water has two of each.[^af2e-10-1] That is the pair's contrast in a single row: H₂O is a liquid and H₂S a gas because the hydrogen bond, not the heavier mass, sets the boiling point. ## Effect on the behavior of gases The [[Kinetic_theory_of_gases|kinetic theory]] behind the [[Ideal_gas_law|ideal gas law]] assumes, among its postulates, that molecules exert no forces on one another between collisions.[^averill-kmt] Real gases obey PV = nRT closely only at low [[Pressure|pressure]] and high temperature, where the molecules are far apart and moving fast enough for the attractions to be negligible.[^averill-real] Van der Waals's equation, `(P + a·n²/V²)·(V − n·b) = n·R·T`, restores the two missing effects: the term in a lowers the measured pressure because molecules near the wall are pulled back by their neighbours, and b subtracts the volume the molecules themselves occupy.[^up2-vdw] The constant a is a direct measure of intermolecular attraction, and it ranks gases the way the microsim's ladders do, smallest for helium and growing through nitrogen and carbon dioxide to ammonia and water, whose hydrogen bonding makes them the most non-ideal of the common gases.[^averill-real] The same attractions decide whether a gas can be liquefied by pressure at all. Above its critical temperature no pressure will condense a gas, because the kinetic energy of the molecules outruns the attractive potential; the critical temperature is 647 K for water and only 5.3 K for helium, so helium stays a gas at room temperature under any pressure while steam at 100 °C condenses under one atmosphere.[^af2e-crit][^up2-crit] Below the [[Critical_point_(thermodynamics)|critical point]], the [[Vapor_pressure|vapour pressure]] of a liquid at a given temperature is likewise a report on its intermolecular forces: the weaker the attractions, the more molecules escape and the higher the pressure, which is why methane's vapour pressure reaches one atmosphere at 112 K and water's only at 373 K.[^nist-webbook] ## Quantum mechanical theories Classical electrostatics accounts for the Keesom and Debye terms and for the ion–dipole force, but the dispersion force has no classical explanation, since a molecule with no permanent dipole should not attract another. London's 1930 treatment derived it as a second-order perturbation of the [[Schrödinger_equation|Schrödinger equation]] for two coupled molecules: the correlated [[Zero-point_energy|zero-point motion]] of the electrons lowers the energy of the pair, and the r⁻⁶ law and the dependence on polarizability and ionization energy follow from the perturbation sum.[^london1930][^london1937] Casimir and Polder showed in 1948 that at separations comparable to the wavelength of the electronic transitions the finite speed of light weakens the correlation, so the dispersion energy falls as r⁻⁷ instead of r⁻⁶ at long range.[^casimir1948] Modern [[Quantum_chemistry|quantum chemistry]] computes intermolecular energies in two ways. Symmetry-adapted perturbation theory treats the interaction between two molecules as a perturbation and returns the energy already partitioned into electrostatic, exchange-repulsion, induction and dispersion terms, so a hydrogen bond or a stacked pair of aromatic rings can be given a physical breakdown.[^sapt1994] Supermolecular methods instead compute the pair and the separated molecules and subtract; standard [[Density_functional_theory|density functional theory]] misses dispersion almost entirely, and Grimme's DFT-D corrections, parametrised in 2010 for the elements from hydrogen to plutonium, add an empirical r⁻⁶ series that restores it.[^grimme2010] These are the potentials that [[Molecular_dynamics|molecular dynamics]] simulations of liquids, [[Polymer|polymers]] and proteins integrate; the hydride ladder in the microsim is a coarse experimental check on the same quantum sums. ## See also - [[Van_der_Waals_force]] - [[London_dispersion_force]] - [[Non-covalent_interaction]] - [[Dipole]] - [[Hydrogen_bond]] - [[Chemical_bond]] - [[Vapor_pressure]] - [[Boiling_point]] ## References [^iupac-hb]: Arunan, E.; Desiraju, G. R.; Klein, R. A.; et al. (2011). "Definition of the hydrogen bond (IUPAC Recommendations 2011)." *Pure and Applied Chemistry* 83 (8): 1637–1641. https://doi.org/10.1351/PAC-REC-10-01-02 [^af2e-10-1]: Flowers, Paul; Neth, Edward; Robinson, William; et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 10 "Liquids and Solids", §10.1 "Intermolecular Forces" (chapter pp. 475–544; page to pin). https://openstax.org/books/chemistry-atoms-first-2e/pages/10-1-intermolecular-forces — Portal Book 051, https://open.umn.edu/opentextbooks/textbooks/chemistry-atoms-first [^af2e-crit]: Flowers et al. (2019), *Chemistry: Atoms First 2e*, Chapter 10, §10.4 "Phase Diagrams", pp. 506–508 (critical temperature and liquefaction). Portal Book 051. [^nist-webbook]: NIST Chemistry WebBook, SRD 69. Phase-change data (normal boiling point, enthalpy of vaporization) for each species, looked up by name or CAS number (water 7732-18-5; hydrogen sulfide 7783-06-4; ammonia 7664-41-7; hydrogen fluoride 7664-39-3; methane 74-82-8). https://webbook.nist.gov/chemistry/ [^cboc-ch6]: Blackstock, Lindsay; Brewer, Sharon; Cinel, Bruno (2022). *Chemical Bonding and Organic Chemistry*. Chapter 6 "Intermolecular Forces and Liquids and Solids" (chapter pp. 363–381; page to pin). Portal Book 054, https://open.umn.edu/opentextbooks/textbooks/chemical-bonding-and-organic-chemistry [^cboc-h2]: Blackstock, Brewer and Cinel (2022), *Chemical Bonding and Organic Chemistry*, Chapter 4 "Chemical Bonding I", p. 231 (H₂ → 2H, +436 kJ/mol). Portal Book 054. [^cboc-nacl]: Blackstock, Brewer and Cinel (2022), *Chemical Bonding and Organic Chemistry*, Chapter 4 "Chemical Bonding I", pp. 224–225 (Born–Haber cycle for NaCl; lattice enthalpy −787 kJ/mol). Portal Book 054. [^averill-kmt]: Averill, Bruce; Eldredge, Patricia (2011). *General Chemistry: Principles, Patterns, and Applications*. Chapter 10 "Gases", pp. 938–939 (postulates of the kinetic molecular theory). Portal Book 050, https://open.umn.edu/opentextbooks/textbooks/general-chemistry-principles-patterns-and-applications [^averill-real]: Averill and Eldredge (2011), *General Chemistry*, Chapter 10 "Gases", pp. 954–976 (real gases and the van der Waals constants; page to pin). Portal Book 050. [^up2-vdw]: Sanny, Jeff; Ling, Samuel J.; et al. (2016). *University Physics Volume 2*. OpenStax. Chapter 2 "The Kinetic Theory of Gases", p. 83 (the van der Waals equation of state). Portal Book 078. [^up2-crit]: Sanny, Ling et al. (2016), *University Physics Volume 2*, Chapter 2, Table 2.1, p. 85 (critical temperatures and pressures: water 647.4 K, helium 5.3 K). Portal Book 078. [^watson-crick]: Watson, J. D.; Crick, F. H. C. (1953). "Molecular Structure of Nucleic Acids: A Structure for Deoxyribose Nucleic Acid." *Nature* 171: 737–738. https://doi.org/10.1038/171737a0 [^barlow1983]: Barlow, D. J.; Thornton, J. M. (1983). "Ion-pairs in proteins." *Journal of Molecular Biology* 168 (4): 867–885. [^kumar2002]: Kumar, Sandeep; Nussinov, Ruth (2002). "Close-range electrostatic interactions in proteins." *ChemBioChem* 3 (7): 604–617. [^vdw-nobel]: Nobel Prize Outreach. "The Nobel Prize in Physics 1910 — Johannes Diderik van der Waals." https://www.nobelprize.org/prizes/physics/1910/summary/ [^lj1924]: Lennard-Jones, J. E. (1924). "On the determination of molecular fields. II. From the equation of state of a gas." *Proceedings of the Royal Society A* 106 (738): 463–477. [^keesom1921]: Keesom, W. H. (1921). "Van der Waals attractive force." *Physikalische Zeitschrift* 22: 129–141. [^debye1920]: Debye, P. (1920). "Die van der Waalsschen Kohäsionskräfte." *Physikalische Zeitschrift* 21: 178–187. [^london1930]: London, F. (1930). "Zur Theorie und Systematik der Molekularkräfte." *Zeitschrift für Physik* 63: 245–279. [^london1937]: London, F. (1937). "The general theory of molecular forces." *Transactions of the Faraday Society* 33: 8–26. [^autumn2002]: Autumn, Kellar; Sitti, Metin; Liang, Yiching A.; et al. (2002). "Evidence for van der Waals adhesion in gecko setae." *Proceedings of the National Academy of Sciences* 99 (19): 12252–12256. [^casimir1948]: Casimir, H. B. G.; Polder, D. (1948). "The Influence of Retardation on the London–van der Waals Forces." *Physical Review* 73 (4): 360–372. https://doi.org/10.1103/PhysRev.73.360 [^sapt1994]: Jeziorski, Bogumił; Moszynski, Robert; Szalewicz, Krzysztof (1994). "Perturbation Theory Approach to Intermolecular Potential Energy Surfaces of van der Waals Complexes." *Chemical Reviews* 94 (7): 1887–1930. [^grimme2010]: Grimme, Stefan; Antony, Jens; Ehrlich, Stephan; Krieg, Helge (2010). "A consistent and accurate ab initio parametrization of density functional dispersion correction (DFT-D) for the 94 elements H-Pu." *Journal of Chemical Physics* 132 (15): 154104. https://doi.org/10.1063/1.3382344 <!-- MATTERSIM:BEGIN g33 — Matter & Energy Cluster microsim (framework build, specs/sims/Intermolecular_force.json); do not hand-edit inside --> **Microsim — three.js (Wikitube framework):** *Intermolecular force* <div class="wt-sim" data-src="https://wikitube-3d-microsims.netlify.app/matter/Intermolecular_force.html" data-title="Intermolecular force"></div> *Built from `MICROSIM_GUIDE/specs/sims/Intermolecular_force.json`; part of the [[PORTAL_Matter|Matter portal]] spine (section sims and See-also variants).* <!-- MATTERSIM:END --> ## Wikipedia : Wikitube **Strict pair:** [Wikipedia](https://en.wikipedia.org/wiki/Intermolecular_force) : [Wikitube](https://en.wikitube.io/wiki/Intermolecular_force) · pinned revision [1366432922](https://en.wikipedia.org/w/index.php?oldid=1366432922) · 2026-09-11 ## Previous hub tags Hubs: `Life_Physics`. Portals: [[PORTAL_Chemistry]]. --- *Matter & Energy Cluster child articles, wave 1 · 2026-09-11 · drafted · Chemistry row K21 · sim pending (matter/Intermolecular_force).*