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An intermolecular force (IMF; also secondary force) is the force that mediates interaction between molecules, including the electromagnetic forces of attraction or repulsion which act between atoms and other types of neighbouring particles, e.g. atoms or ions. Intermolecular forces are weak relative to intramolecular forces – the forces which hold a molecule together. For example, the covalent bond, involving sharing electron pairs between atoms, is much stronger than the forces present between neighboring molecules.^[1] Both sets of forces are essential parts of force fields frequently used in molecular mechanics.

The first reference to the nature of microscopic forces is found in Alexis Clairaut's work Théorie de la figure de la Terre, published in Paris in 1743.^[2] Other scientists who have contributed to the investigation of microscopic forces include: Laplace, Gauss, Maxwell, Boltzmann and Pauling.

Attractive intermolecular forces are categorized into the following types:

Hydrogen bonding
Ion–dipole forces and ion–induced dipole force
Cation–π, σ–π and π–π bonding
Van der Waals forces – Keesom force, Debye force, and London dispersion force
Cation–cation bonding
Salt bridge (protein and supramolecular)

Information on intermolecular forces is obtained by macroscopic measurements of properties like viscosity, pressure, volume, temperature (PVT) data. The link to microscopic aspects is given by virial coefficients and intermolecular pair potentials, such as the Mie potential, Buckingham potential or Lennard-Jones potential.

In the broadest sense, it can be understood as such interactions between any particles (molecules, atoms, ions and molecular ions) in which the formation of chemical (that is, ionic, covalent or metallic) bonds does not occur. In other words, these interactions are significantly weaker than covalent ones and do not lead to a significant restructuring of the electronic structure of the interacting particles. (This is only partially true. For example, all enzymatic and catalytic reactions begin with a weak intermolecular interaction between a substrate and an enzyme or a molecule with a catalyst, but several such weak interactions with the required spatial configuration of the active center of the enzyme lead to significant restructuring changes the energy state of molecules or substrate, which ultimately leads to the breaking of some and the formation of other covalent chemical bonds. Strictly speaking, all enzymatic reactions begin with intermolecular interactions between the substrate and the enzyme, therefore the importance of these interactions is especially great in biochemistry and molecular biology,^[3] and is the basis of enzymology).

Hydrogen bonding

A hydrogen bond is an extreme form of dipole-dipole bonding, referring to the attraction between a hydrogen atom that is bonded to an element with high electronegativity, usually nitrogen, oxygen, or fluorine.^[4] The hydrogen bond is often described as a strong electrostatic dipole–dipole interaction. However, it also has some features of covalent bonding: it is directional, stronger than a van der Waals force interaction, produces interatomic distances shorter than the sum of their van der Waals radii, and usually involves a limited number of interaction partners, which can be interpreted as a kind of valence. The number of Hydrogen bonds formed between molecules is equal to the number of active pairs. The molecule which donates its hydrogen is termed the donor molecule, while the molecule containing lone pair participating in H bonding is termed the acceptor molecule. The number of active pairs is equal to the common number between number of hydrogens the donor has and the number of lone pairs the acceptor has.

Though both not depicted in the diagram, water molecules have four active bonds. The oxygen atom’s two lone pairs interact with a hydrogen each, forming two additional hydrogen bonds, and the second hydrogen atom also interacts with a neighbouring oxygen. Intermolecular hydrogen bonding is responsible for the high boiling point of water (100 °C) compared to the other group 16 hydrides, which have little capability to hydrogen bond. Intramolecular hydrogen bonding is partly responsible for the secondary, tertiary, and quaternary structures of proteins and nucleic acids. It also plays an important role in the structure of polymers, both synthetic and natural.^[5]

Salt bridge

The attraction between cationic and anionic sites is a noncovalent, or intermolecular interaction which is usually referred to as ion pairing or salt bridge.^[6] It is essentially due to electrostatic forces, although in aqueous medium the association is driven by entropy and often even endothermic. Most salts form crystals with characteristic distances between the ions; in contrast to many other noncovalent interactions, salt bridges are not directional and show in the solid state usually contact determined only by the van der Waals radii of the ions. Inorganic as well as organic ions display in water at moderate ionic strength I similar salt bridge as association ΔG values around 5 to 6 kJ/mol for a 1:1 combination of anion and cation, almost independent of the nature (size, polarizability, etc.) of the ions.^[7] The ΔG values are additive and approximately a linear function of the charges, the interaction of e.g. a doubly charged phosphate anion with a single charged ammonium cation accounts for about 2x5 = 10 kJ/mol. The ΔG values depend on the ionic strength I of the solution, as described by the Debye-Hückel equation, at zero ionic strength one observes ΔG = 8 kJ/mol.

Dipole–dipole and similar interactions

Dipole–dipole interactions (or Keesom interactions) are electrostatic interactions between molecules which have permanent dipoles. This interaction is stronger than the London forces but is weaker than ion-ion interaction because only partial charges are involved. These interactions tend to align the molecules to increase attraction (reducing potential energy). An example of a dipole–dipole interaction can be seen in hydrogen chloride (HCl): the positive end of a polar molecule will attract the negative end of the other molecule and influence its position. Polar molecules have a net attraction between them. Examples of polar molecules include hydrogen chloride (HCl) and chloroform (CHCl₃).

{\overset {\color {Red}\delta +}{{\ce {H}}}}-{\overset {\color {Red}\delta -}{{\ce {Cl}}}}\cdots {\overset {\color {Red}\delta +}{{\ce {H}}}}-{\overset {\color {Red}\delta -}{{\ce {Cl}}}}

Often molecules contain dipolar groups of atoms, but have no overall dipole moment on the molecule as a whole. This occurs if there is symmetry within the molecule that causes the dipoles to cancel each other out. This occurs in molecules such as tetrachloromethane and carbon dioxide. The dipole–dipole interaction between two individual atoms is usually zero, since atoms rarely carry a permanent dipole.

The Keesom interaction is a van der Waals force. It is discussed further in the section "Van der Waals forces".

Ion–dipole and ion–induced dipole forces

Ion–dipole and ion–induced dipole forces are similar to dipole–dipole and dipole–induced dipole interactions but involve ions, instead of only polar and non-polar molecules. Ion–dipole and ion–induced dipole forces are stronger than dipole–dipole interactions because the charge of any ion is much greater than the charge of a dipole moment. Ion–dipole bonding is stronger than hydrogen bonding.^[8]

An ion–dipole force consists of an ion and a polar molecule interacting. They align so that the positive and negative groups are next to one another, allowing maximum attraction. An important example of this interaction is hydration of ions in water which give rise to hydration enthalpy. The polar water molecules surround themselves around ions in water and the energy released during the process is known as hydration enthalpy. The interaction has its immense importance in justifying the stability of various ions (like Cu²⁺) in water.

An ion–induced dipole force consists of an ion and a non-polar molecule interacting. Like a dipole–induced dipole force, the charge of the ion causes distortion of the electron cloud on the non-polar molecule.^[9]

Van der Waals forces

The van der Waals forces arise from interaction between uncharged atoms or molecules, leading not only to such phenomena as the cohesion of condensed phases and physical absorption of gases, but also to a universal force of attraction between macroscopic bodies.^[10]

Keesom force (permanent dipole – permanent dipole)

The first contribution to van der Waals forces is due to electrostatic interactions between rotating permanent dipoles, quadrupoles (all molecules with symmetry lower than cubic), and multipoles. It is termed the Keesom interaction, named after Willem Hendrik Keesom.^[11] These forces originate from the attraction between permanent dipoles (dipolar molecules) and are temperature dependent.^[10]

They consist of attractive interactions between dipoles that are ensemble averaged over different rotational orientations of the dipoles. It is assumed that the molecules are constantly rotating and never get locked into place. This is a good assumption, but at some point molecules do get locked into place. The energy of a Keesom interaction depends on the inverse sixth power of the distance, unlike the interaction energy of two spatially fixed dipoles, which depends on the inverse third power of the distance. The Keesom interaction can only occur among molecules that possess permanent dipole moments, i.e., two polar molecules. Also Keesom interactions are very weak van der Waals interactions and do not occur in aqueous solutions that contain electrolytes. The angle averaged interaction is given by the following equation:

{\frac {-d_{1}^{2}d_{2}^{2}}{24\pi ^{2}\varepsilon _{0}^{2}\varepsilon _{r}^{2}k_{\text{B}}Tr^{6}}}=V,

where d = electric dipole moment, $\varepsilon _{0}$ = permittivity of free space, $\varepsilon _{r}$ = dielectric constant of surrounding material, T = temperature, $k_{\text{B}}$ = Boltzmann constant, and r = distance between molecules.

Debye force (permanent dipoles–induced dipoles)

The second contribution is the induction (also termed polarization) or Debye force, arising from interactions between rotating permanent dipoles and from the polarizability of atoms and molecules (induced dipoles). These induced dipoles occur when one molecule with a permanent dipole repels another molecule's electrons. A molecule with permanent dipole can induce a dipole in a similar neighboring molecule and cause mutual attraction. Debye forces cannot occur between atoms. The forces between induced and permanent dipoles are not as temperature dependent as Keesom interactions because the induced dipole is free to shift and rotate around the polar molecule. The Debye induction effects and Keesom orientation effects are termed polar interactions.^[10]

The induced dipole forces appear from the induction (also termed polarization), which is the attractive interaction between a permanent multipole on one molecule with an induced (by the former di/multi-pole) 31 on another.^[12]^[13]^[14] This interaction is called the Debye force, named after Peter J. W. Debye.

One example of an induction interaction between permanent dipole and induced dipole is the interaction between HCl and Ar. In this system, Ar experiences a dipole as its electrons are attracted (to the H side of HCl) or repelled (from the Cl side) by HCl.^[12]^[13] The angle averaged interaction is given by the following equation:

{\frac {-d_{1}^{2}\alpha _{2}}{16\pi ^{2}\varepsilon _{0}^{2}\varepsilon _{r}^{2}r^{6}}}=V,

where $\alpha _{2}$ = polarizability.

This kind of interaction can be expected between any polar molecule and non-polar/symmetrical molecule. The induction-interaction force is far weaker than dipole–dipole interaction, but stronger than the London dispersion force.

London dispersion force (fluctuating dipole–induced dipole interaction)

The third and dominant contribution is the dispersion or London force (fluctuating dipole–induced dipole), which arises due to the non-zero instantaneous dipole moments of all atoms and molecules. Such polarization can be induced either by a polar molecule or by the repulsion of negatively charged electron clouds in non-polar molecules. Thus, London interactions are caused by random fluctuations of electron density in an electron cloud. An atom with a large number of electrons will have a greater associated London force than an atom with fewer electrons. The dispersion (London) force is the most important component because all materials are polarizable, whereas Keesom and Debye forces require permanent dipoles. The London interaction is universal and is present in atom-atom interactions as well. For various reasons, London interactions (dispersion) have been considered relevant for interactions between macroscopic bodies in condensed systems. Hamaker developed the theory of van der Waals between macroscopic bodies in 1937 and showed that the additivity of these interactions renders them considerably more long-range.^[10]

Relative strength of forces

Bond type	Dissociation energy (kcal/mol)^[15]	Dissociation energy (kJ/mol)	Note
Ionic lattice	250–4000^[16]	1100–20000
Covalent bond	30–260	130–1100
Hydrogen bond	1–12	4–50	About 5 kcal/mol (21 kJ/mol) in water
Dipole–dipole	0.5–2	2–8
London dispersion forces	<1 to 15	<4 to 63	Estimated from the enthalpies of vaporization of hydrocarbons^[17]

This comparison is approximate. The actual relative strengths will vary depending on the molecules involved. For instance, the presence of water creates competing interactions that greatly weaken the strength of both ionic and hydrogen bonds.^[18] We may consider that for static systems, Ionic bonding and covalent bonding will always be stronger than intermolecular forces in any given substance. But it is not so for big moving systems like enzyme molecules interacting with substrate molecules.^[19] Here the numerous intramolecular (most often - hydrogen bonds) bonds form an active intermediate state where the intermolecular bonds cause some of the covalent bond to be broken, while the others are formed, in this way proceeding the thousands of enzymatic reactions, so important for living organisms.

Effect on the behavior of gases

Intermolecular forces are repulsive at short distances and attractive at long distances (see the Lennard-Jones potential).^[20]^[21] In a gas, the repulsive force chiefly has the effect of keeping two molecules from occupying the same volume. This gives a real gas a tendency to occupy a larger volume than an ideal gas at the same temperature and pressure. The attractive force draws molecules closer together and gives a real gas a tendency to occupy a smaller volume than an ideal gas. Which interaction is more important depends on temperature and pressure (see compressibility factor).

In a gas, the distances between molecules are generally large, so intermolecular forces have only a small effect. The attractive force is not overcome by the repulsive force, but by the thermal energy of the molecules. Temperature is the measure of thermal energy, so increasing temperature reduces the influence of the attractive force. In contrast, the influence of the repulsive force is essentially unaffected by temperature.

When a gas is compressed to increase its density, the influence of the attractive force increases. If the gas is made sufficiently dense, the attractions can become large enough to overcome the tendency of thermal motion to cause the molecules to disperse. Then the gas can condense to form a solid or liquid, i.e., a condensed phase. Lower temperature favors the formation of a condensed phase. In a condensed phase, there is very nearly a balance between the attractive and repulsive forces.

Quantum mechanical theories

Intermolecular forces observed between atoms and molecules can be described phenomenologically as occurring between permanent and instantaneous dipoles, as outlined above. Alternatively, one may seek a fundamental, unifying theory that is able to explain the various types of interactions such as hydrogen bonding,^[22] van der Waals force^[23] and dipole–dipole interactions. Typically, this is done by applying the ideas of quantum mechanics to molecules, and Rayleigh–Schrödinger perturbation theory has been especially effective in this regard. When applied to existing quantum chemistry methods, such a quantum mechanical explanation of intermolecular interactions provides an array of approximate methods that can be used to analyze intermolecular interactions.^[24] One of the most helpful methods to visualize this kind of intermolecular interactions, that we can find in quantum chemistry, is the non-covalent interaction index, which is based on the electron density of the system. London dispersion forces play a big role with this.

Concerning electron density topology, recent methods based on electron density gradient methods have emerged recently, notably with the development of IBSI (Intrinsic Bond Strength Index),^[25] relying on the IGM (Independent Gradient Model) methodology.^[26]^[27]^[28]

References

^ Fischer, Johann; Wendland, Martin (October 2023). "On the history of key empirical intermolecular potentials". Fluid Phase Equilibria. 573: 113876. Bibcode:2023FlPEq.57313876F. doi:10.1016/j.fluid.2023.113876.
^ Margenau H, Kestner NR (1969). Theory of Intermolecular Forces. International Series of Monographs in Natural Philosophy. Vol. 18 (1st ed.). Oxford: Pergamon Press. ISBN 978-0-08-016502-8.
^ "Biochemistry and Molecular Biology - Paperback - Despo Papachristodoulou, Alison Snape, William H. Elliott, Daphne C. Elliott - Oxford University Press". global.oup.com. Retrieved 2024-01-04.
^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "hydrogen bond". doi:10.1351/goldbook.H02899
^ Lindh U (2013), "Biological functions of the elements", in Selinus O (ed.), Essentials of Medical Geology (Revised ed.), Dordrecht: Springer, pp. 129–177, doi:10.1007/978-94-007-4375-5_7, ISBN 978-94-007-4374-8
^ Ciferri A, Perico A, eds. (2012). Ionic Interactions in Natural and Synthetic Macromolecules. Hoboken, NJ: John Wiley & Sons, Inc. ISBN 978-0-470-52927-0.
^ Biedermann F, Schneider HJ (May 2016). "Experimental Binding Energies in Supramolecular Complexes". Chemical Reviews. 116 (9): 5216–5300. doi:10.1021/acs.chemrev.5b00583. PMID 27136957.
^ Tro N (2011). Chemistry: A Molecular Approach. United States: Pearson Education Inc. p. 466. ISBN 978-0-321-65178-5.
^ Blaber M (1996). "Intermolecular Forces". mikeblaber.org. Archived from the original on 2020-08-01. Retrieved 2011-11-17.
^ ^a ^b ^c ^d Leite FL, Bueno CC, Da Róz AL, Ziemath EC, Oliveira ON (October 2012). "Theoretical models for surface forces and adhesion and their measurement using atomic force microscopy". International Journal of Molecular Sciences. 13 (10): 12773–12856. doi:10.3390/ijms131012773. PMC 3497299. PMID 23202925.
^ Keesom WH (1915). "The second virial coefficient for rigid spherical molecules whose mutual attraction is equivalent to that of a quadruplet placed at its center" (PDF). Proceedings of the Royal Netherlands Academy of Arts and Sciences. 18: 636–646.
^ ^a ^b Blustin PH (1978). "A Floating Gaussian Orbital calculation on argon hydrochloride (Ar·HCl)". Theoretica Chimica Acta. 47 (3): 249–257. doi:10.1007/BF00577166. S2CID 93104668.
^ ^a ^b Roberts JK, Orr WJ (1938). "Induced dipoles and the heat of adsorption of argon on ionic crystals". Transactions of the Faraday Society. 34: 1346. doi:10.1039/TF9383401346.
^ Sapse AM, Rayez-Meaume MT, Rayez JC, Massa LJ (1979). "Ion-induced dipole H−n clusters". Nature. 278 (5702): 332–333. Bibcode:1979Natur.278..332S. doi:10.1038/278332a0. S2CID 4304250.
^ Eğe SN (2004). Organic Chemistry: Structure and Reactivity (5th ed.). Boston: Houghton Mifflin Company. pp. 30–33, 67. ISBN 978-0-618-31809-4.
^ "Lattice Energies". Division of Chemical Education. Purdue University. Retrieved 2014-01-21.
^ Majer V, Svoboda V (1985). Enthalpies of Vaporization of Organic Compounds. Oxford: Blackwell Scientific. ISBN 978-0-632-01529-0.
^ Alberts, Bruce; et al. (2015). Molecular biology of the cell (6th ed.). New York, NY: Garland Science - Taylor & Francis. ISBN 978-0-8153-4432-2. OCLC 887605755.
^ Savir Y, Tlusty T (May 2007). "Conformational proofreading: the impact of conformational changes on the specificity of molecular recognition". PLOS ONE. 2 (5): e468. Bibcode:2007PLoSO...2..468S. doi:10.1371/journal.pone.0000468. PMC 1868595. PMID 17520027.
^ Fischer, Johann; Wendland, Martin (October 2023). "On the history of key empirical intermolecular potentials". Fluid Phase Equilibria. 573: 113876. Bibcode:2023FlPEq.57313876F. doi:10.1016/j.fluid.2023.113876.
^ Lenhard, Johannes; Stephan, Simon; Hasse, Hans (June 2024). "On the History of the Lennard-Jones Potential". Annalen der Physik. 536 (6). doi:10.1002/andp.202400115. ISSN 0003-3804.
^ Arunan E, Desiraju GR, Klein RA, Sadlej J, Scheiner S, Alkorta I, et al. (2011-07-08). "Definition of the hydrogen bond (IUPAC Recommendations 2011)". Pure and Applied Chemistry. 83 (8): 1637–1641. doi:10.1351/PAC-REC-10-01-02. ISSN 1365-3075. S2CID 97688573.
^ Landau LD, Lifshitz EM (1960). Electrodynamics of Continuous Media. Oxford: Pergamon. pp. 368–376.
^ King M (1976). "Theory of the Chemical Bond". Journal of the American Chemical Society. 98 (12): 3415–3420. doi:10.1021/ja00428a004.
^ Klein J, Khartabil H, Boisson JC, Contreras-García J, Piquemal JP, Hénon E (March 2020). "New Way for Probing Bond Strength" (PDF). The Journal of Physical Chemistry A. 124 (9): 1850–1860. Bibcode:2020JPCA..124.1850K. doi:10.1021/acs.jpca.9b09845. PMID 32039597. S2CID 211070812.
^ Lefebvre C, Rubez G, Khartabil H, Boisson JC, Contreras-García J, Hénon E (July 2017). "Accurately extracting the signature of intermolecular interactions present in the NCI plot of the reduced density gradient versus electron density" (PDF). Physical Chemistry Chemical Physics. 19 (27): 17928–17936. Bibcode:2017PCCP...1917928L. doi:10.1039/C7CP02110K. PMID 28664951.
^ Lefebvre C, Khartabil H, Boisson JC, Contreras-García J, Piquemal JP, Hénon E (March 2018). "The Independent Gradient Model: A New Approach for Probing Strong and Weak Interactions in Molecules from Wave Function Calculations" (PDF). ChemPhysChem. 19 (6): 724–735. doi:10.1002/cphc.201701325. PMID 29250908.
^ Ponce-Vargas M, Lefebvre C, Boisson JC, Hénon E (January 2020). "Atomic Decomposition Scheme of Noncovalent Interactions Applied to Host-Guest Assemblies". Journal of Chemical Information and Modeling. 60 (1): 268–278. doi:10.1021/acs.jcim.9b01016. PMID 31877034. S2CID 209488458.

[1] Fischer, Johann; Wendland, Martin (October 2023). "On the history of key empirical intermolecular potentials". Fluid Phase Equilibria. 573: 113876. Bibcode:2023FlPEq.57313876F. doi:10.1016/j.fluid.2023.113876.

[2] Margenau H, Kestner NR (1969). Theory of Intermolecular Forces. International Series of Monographs in Natural Philosophy. Vol. 18 (1st ed.). Oxford: Pergamon Press. ISBN 978-0-08-016502-8.

[3] "Biochemistry and Molecular Biology - Paperback - Despo Papachristodoulou, Alison Snape, William H. Elliott, Daphne C. Elliott - Oxford University Press". global.oup.com. Retrieved 2024-01-04.

[GoldBook-H02899-4] IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "hydrogen bond". doi:10.1351/goldbook.H02899

[5] Lindh U (2013), "Biological functions of the elements", in Selinus O (ed.), Essentials of Medical Geology (Revised ed.), Dordrecht: Springer, pp. 129–177, doi:10.1007/978-94-007-4375-5_7, ISBN 978-94-007-4374-8

[6] Ciferri A, Perico A, eds. (2012). Ionic Interactions in Natural and Synthetic Macromolecules. Hoboken, NJ: John Wiley & Sons, Inc. ISBN 978-0-470-52927-0.

[7] Biedermann F, Schneider HJ (May 2016). "Experimental Binding Energies in Supramolecular Complexes". Chemical Reviews. 116 (9): 5216–5300. doi:10.1021/acs.chemrev.5b00583. PMID 27136957.

[8] Tro N (2011). Chemistry: A Molecular Approach. United States: Pearson Education Inc. p. 466. ISBN 978-0-321-65178-5.

[Michael-Blaber-1996-9] Blaber M (1996). "Intermolecular Forces". mikeblaber.org. Archived from the original on 2020-08-01. Retrieved 2011-11-17.

[j1-10] Leite FL, Bueno CC, Da Róz AL, Ziemath EC, Oliveira ON (October 2012). "Theoretical models for surface forces and adhesion and their measurement using atomic force microscopy". International Journal of Molecular Sciences. 13 (10): 12773–12856. doi:10.3390/ijms131012773. PMC 3497299. PMID 23202925.

[11] Keesom WH (1915). "The second virial coefficient for rigid spherical molecules whose mutual attraction is equivalent to that of a quadruplet placed at its center" (PDF). Proceedings of the Royal Netherlands Academy of Arts and Sciences. 18: 636–646.

[Blustin-1978-12] Blustin PH (1978). "A Floating Gaussian Orbital calculation on argon hydrochloride (Ar·HCl)". Theoretica Chimica Acta. 47 (3): 249–257. doi:10.1007/BF00577166. S2CID 93104668.

[Roberts-Orr-1938-13] Roberts JK, Orr WJ (1938). "Induced dipoles and the heat of adsorption of argon on ionic crystals". Transactions of the Faraday Society. 34: 1346. doi:10.1039/TF9383401346.

[Sapse-et-al-1979-14] Sapse AM, Rayez-Meaume MT, Rayez JC, Massa LJ (1979). "Ion-induced dipole H−n clusters". Nature. 278 (5702): 332–333. Bibcode:1979Natur.278..332S. doi:10.1038/278332a0. S2CID 4304250.

[Seyhan-Organic-Chemistry-15] Eğe SN (2004). Organic Chemistry: Structure and Reactivity (5th ed.). Boston: Houghton Mifflin Company. pp. 30–33, 67. ISBN 978-0-618-31809-4.

[Purdue-Lattice-16] "Lattice Energies". Division of Chemical Education. Purdue University. Retrieved 2014-01-21.

[Majer-Svoboda-enthalpy-vap-17] Majer V, Svoboda V (1985). Enthalpies of Vaporization of Organic Compounds. Oxford: Blackwell Scientific. ISBN 978-0-632-01529-0.

[18] Alberts, Bruce; et al. (2015). Molecular biology of the cell (6th ed.). New York, NY: Garland Science - Taylor & Francis. ISBN 978-0-8153-4432-2. OCLC 887605755.

[19] Savir Y, Tlusty T (May 2007). "Conformational proofreading: the impact of conformational changes on the specificity of molecular recognition". PLOS ONE. 2 (5): e468. Bibcode:2007PLoSO...2..468S. doi:10.1371/journal.pone.0000468. PMC 1868595. PMID 17520027.

[20] Fischer, Johann; Wendland, Martin (October 2023). "On the history of key empirical intermolecular potentials". Fluid Phase Equilibria. 573: 113876. Bibcode:2023FlPEq.57313876F. doi:10.1016/j.fluid.2023.113876.

[21] Lenhard, Johannes; Stephan, Simon; Hasse, Hans (June 2024). "On the History of the Lennard-Jones Potential". Annalen der Physik. 536 (6). doi:10.1002/andp.202400115. ISSN 0003-3804.

[:0-22] Arunan E, Desiraju GR, Klein RA, Sadlej J, Scheiner S, Alkorta I, et al. (2011-07-08). "Definition of the hydrogen bond (IUPAC Recommendations 2011)". Pure and Applied Chemistry. 83 (8): 1637–1641. doi:10.1351/PAC-REC-10-01-02. ISSN 1365-3075. S2CID 97688573.

[LD-23] Landau LD, Lifshitz EM (1960). Electrodynamics of Continuous Media. Oxford: Pergamon. pp. 368–376.

[24] King M (1976). "Theory of the Chemical Bond". Journal of the American Chemical Society. 98 (12): 3415–3420. doi:10.1021/ja00428a004.

[25] Klein J, Khartabil H, Boisson JC, Contreras-García J, Piquemal JP, Hénon E (March 2020). "New Way for Probing Bond Strength" (PDF). The Journal of Physical Chemistry A. 124 (9): 1850–1860. Bibcode:2020JPCA..124.1850K. doi:10.1021/acs.jpca.9b09845. PMID 32039597. S2CID 211070812.

[26] Lefebvre C, Rubez G, Khartabil H, Boisson JC, Contreras-García J, Hénon E (July 2017). "Accurately extracting the signature of intermolecular interactions present in the NCI plot of the reduced density gradient versus electron density" (PDF). Physical Chemistry Chemical Physics. 19 (27): 17928–17936. Bibcode:2017PCCP...1917928L. doi:10.1039/C7CP02110K. PMID 28664951.

[27] Lefebvre C, Khartabil H, Boisson JC, Contreras-García J, Piquemal JP, Hénon E (March 2018). "The Independent Gradient Model: A New Approach for Probing Strong and Weak Interactions in Molecules from Wave Function Calculations" (PDF). ChemPhysChem. 19 (6): 724–735. doi:10.1002/cphc.201701325. PMID 29250908.

[28] Ponce-Vargas M, Lefebvre C, Boisson JC, Hénon E (January 2020). "Atomic Decomposition Scheme of Noncovalent Interactions Applied to Host-Guest Assemblies". Journal of Chemical Information and Modeling. 60 (1): 268–278. doi:10.1021/acs.jcim.9b01016. PMID 31877034. S2CID 209488458.

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@@ Line 1: / Line 1: @@
+{{short description|Force of attraction or repulsion between molecules and neighboring particles}}
-[[ja:&#20998;&#23376;&#38291;&#21147;]] [[pl:Oddzialywania miedzyczasteczkowe]]
+An '''intermolecular force''' ('''IMF'''; also '''secondary force''') is the force that mediates interaction between molecules, including the [[Electromagnetism|electromagnetic forces of attraction
-'''Intermolecular forces''' are electromagnetic forces which act between [[molecule|molecules]] or between widely separated regions of a [[macromolecule]]. Listed in order of decreasing strength, these forces are:
+or repulsion]] which act between atoms and other types of neighbouring particles, e.g. [[atom]]s or [[ion]]s. Intermolecular forces are weak relative to [[intramolecular force]]s – the forces which hold a molecule together. For example, the [[covalent bond]], involving sharing electron pairs between atoms, is much stronger than the forces present between neighboring molecules.<ref>{{Cite journal |last1=Fischer |first1=Johann |last2=Wendland |first2=Martin |date=October 2023 |title=On the history of key empirical intermolecular potentials |journal=Fluid Phase Equilibria |language=en |volume=573 |pages=113876 |doi=10.1016/j.fluid.2023.113876|bibcode=2023FlPEq.57313876F |doi-access=free }}</ref> Both sets of forces are essential parts of [[Force field (chemistry)|force fields]] frequently used in [[molecular mechanics]].
+The first reference to the nature of microscopic forces is found in [[Alexis Clairaut]]'s work ''Théorie de la figure de la Terre,'' published in Paris in 1743.<ref>{{cite book | vauthors = Margenau H, Kestner NR | title=Theory of Intermolecular Forces |date=1969 |publisher=Pergamon Press |location=Oxford |isbn=978-0-08-016502-8 |edition=1st | series = International Series of Monographs in Natural Philosophy | volume = 18 }}</ref> Other scientists who have contributed to the investigation of microscopic forces include: [[Pierre-Simon Laplace|Laplace]], [[Carl Friedrich Gauss|Gauss]], [[James Clerk Maxwell|Maxwell]], [[Ludwig Boltzmann|Boltzmann]] and [[Linus Pauling|Pauling]].
-* Ionic interactions
-* [[hydrogen bonding|Hydrogen bonds]]
-* Dipole-dipole interactions
-* [[London forces|London Dispersion Forces]]
+Attractive intermolecular forces are categorized into the following types:
+*[[Hydrogen bond]]ing
+*Ion–dipole forces and ion–induced dipole force
+*[[Cation–π interaction|Cation–π]], σ–π and π–π bonding
+*[[Van der Waals force]]s – [[Keesom force]], [[Debye force]], and [[London dispersion force]]
+*[[Cation–cation bond|Cation–cation bonding]]
+*[[Salt bridge (protein and supramolecular)]]
+Information on intermolecular forces is obtained by macroscopic measurements of properties like [[viscosity]], [[PVT (physics)|pressure, volume, temperature]] (PVT) data. The link to microscopic aspects is given by [[virial coefficient]]s and intermolecular [[Pair potential|pair potentials]], such as the [[Mie potential]], [[Buckingham potential]] or [[Lennard-Jones potential]].
-=== Ionic Interactions ===
+In the broadest sense, it can be understood as such interactions between any particles ([[molecule]]s, [[atom]]s, [[ion]]s and [[molecular ion]]s) in which the formation of chemical (that is, ionic, covalent or metallic) bonds does not occur. In other words, these interactions are significantly weaker than [[Covalent bond|covalent]] ones and do not lead to a significant restructuring of the [[electronic structure]] of the interacting particles. (This is only partially true. For example, all [[Enzyme|enzymatic]] and [[Catalysis|catalytic reactions]] begin with a weak intermolecular interaction between a substrate and an [[enzyme]] or a molecule with a [[Catalyst: Agents of Change|catalyst]], but several such weak interactions with the required spatial configuration of the active center of the enzyme lead to significant restructuring changes the energy state of molecules or substrate, which ultimately leads to the breaking of some and the formation of other covalent chemical bonds. Strictly speaking, all [[Enzyme catalysis|enzymatic reactions]] begin with intermolecular interactions between the [[Substrate (chemistry)|substrate]] and the enzyme, therefore the importance of these interactions is especially great in [[biochemistry]] and [[molecular biology]],<ref>{{Cite web |title=Biochemistry and Molecular Biology - Paperback - Despo Papachristodoulou, Alison Snape, William H. Elliott, Daphne C. Elliott - Oxford University Press |url=https://global.oup.com/ukhe/product/biochemistry-and-molecular-biology-9780198768111 |access-date=2024-01-04 |website=global.oup.com |language=en}}</ref> and is the basis of [[Enzyme|enzymology]]).
-This is an interactions that occurs between charged species, like charges will have a repulsive force while opposite charges will have an attractive force.
-=== Hydrogen Bonding ===
+== Hydrogen bonding ==
+{{Main|Hydrogen bond}}
-[[hydrogen bonding]] can be formed when a hydrogen atom is bound to a highly electronegative atom such as nitrogen, oxygen, or fluorine.  The hydrogen atom has a partial positive charge and can interacts with another electronegative atoms found in a different molecule (again N, O, or F).  This results in a stabilizing interaction that loosely binds the two molecules together.  A common example is [[water]]:
+A ''hydrogen bond'' is an extreme form of dipole-dipole bonding, referring to the attraction between a [[hydrogen]] atom that is bonded to an element with high [[electronegativity]], usually [[nitrogen]], [[oxygen]], or [[fluorine]].<ref name=GoldBook-H02899>{{GoldBookRef |file= H02899 |title= hydrogen bond}}</ref> The hydrogen bond is often described as a strong electrostatic dipole–dipole interaction. However, it also has some features of covalent bonding: it is directional, stronger than a [[van der Waals force]] interaction, produces interatomic distances shorter than the sum of their [[van der Waals radius|van der Waals radii]], and usually involves a limited number of interaction partners, which can be interpreted as a kind of [[Valence (chemistry)|valence]]. The number of Hydrogen bonds formed between molecules is equal to the number of active pairs. The molecule which donates its hydrogen is termed the donor molecule, while the molecule containing lone pair participating in H bonding is termed the acceptor molecule. The number of active pairs is equal to the common number between number of hydrogens the donor has and the number of lone pairs the acceptor has.
-H<sub>2</sub>O---H-O-H
+[[Image:Hydrogen-bonding-in-water-2D.svg|thumb|Hydrogen bonding in water]]
-Hydrogen bonds are found throughout nature.  Hydrogen bonds give water its unique properties, which allowed life to develop on earth; hydrogen bonds are also the intermolecular force that binds together the two strands in a molecule of DNA
+Though both not depicted in the diagram, water molecules have four active bonds. The oxygen atom’s two lone pairs interact with a hydrogen each, forming two additional hydrogen bonds, and the second hydrogen atom also interacts with a neighbouring oxygen. Intermolecular hydrogen bonding is responsible for the high boiling point of [[water]] (100&nbsp;°C) compared to the other [[Hydrogen chalcogenide|group 16 hydrides]], which have little capability to hydrogen bond. Intramolecular hydrogen bonding is partly responsible for the [[secondary structure|secondary]], [[tertiary structure|tertiary]], and [[quaternary structure]]s of [[protein]]s and [[nucleic acid]]s. It also plays an important role in the structure of [[polymers]], both synthetic and natural.<ref>{{Citation | vauthors = Lindh U | contribution = Biological functions of the elements | veditors = Selinus O | title = Essentials of Medical Geology | pages = 129–177 | publisher = Springer | place = Dordrecht | year = 2013 | edition = Revised | isbn = 978-94-007-4374-8 | doi = 10.1007/978-94-007-4375-5_7 }}</ref>
-=== Dipole-Dipole Interactions ===
+==Salt bridge==
-Dipole-Dipole interactions are the forces that occur between two molecules with permanent dipoles.  These work in a similar manner to ionic interactions but are weaker because only partial charges are involved.  An example of this can be seen in [[hydrochloric acid]]:
+{{Main|Salt bridge (protein and supramolecular)}}
+The attraction between cationic and anionic sites is a noncovalent, or intermolecular interaction which is usually referred to as ion pairing or salt bridge.<ref>{{cite book | veditors = Ciferri A, Perico A |title=Ionic Interactions in Natural and Synthetic Macromolecules |date=2012 |publisher=John Wiley & Sons, Inc. |location=Hoboken, NJ |isbn=978-0-470-52927-0}}</ref>
-(+)(-)  &nbsp;&nbsp; (+) (-)<br>
+It is essentially due to electrostatic forces, although in aqueous medium the association is driven by entropy and often even endothermic.  Most salts form crystals with characteristic distances between the ions; in contrast to many other noncovalent interactions, salt bridges are not directional and show in the solid state usually contact determined only by the van der Waals radii of the ions.
-H-Cl-----H-Cl
+Inorganic as well as organic ions display in water at moderate ionic strength I similar salt bridge as association ΔG values around 5 to 6 kJ/mol for a 1:1 combination of anion and cation, almost independent of the nature (size, polarizability, etc.) of the ions.<ref>{{cite journal | vauthors = Biedermann F, Schneider HJ | title = Experimental Binding Energies in Supramolecular Complexes | journal = Chemical Reviews | volume = 116 | issue = 9 | pages = 5216–5300 | date = May 2016 | pmid = 27136957 | doi = 10.1021/acs.chemrev.5b00583 }}</ref> The ΔG values are additive and approximately a linear function of the charges, the interaction of e.g. a doubly charged phosphate anion with a single charged ammonium cation accounts for about 2x5 = 10 kJ/mol. The ΔG  values depend on the ionic strength I of the solution, as described by the Debye-Hückel equation, at zero ionic strength one observes ΔG = 8 kJ/mol.
+==Dipole–dipole and similar interactions {{anchor|Dipole-dipole interactions}}==
-=== London Dispersion Forces ===
+Dipole–dipole interactions (or Keesom interactions) are electrostatic interactions between molecules which have permanent dipoles. This interaction  is stronger than the London forces but is weaker than ion-ion interaction because only partial charges are involved. These interactions tend to align the molecules to increase attraction (reducing [[potential energy]]). An example of a dipole–dipole interaction can be seen in [[hydrogen chloride]] (HCl): the positive end of a polar molecule will attract the negative end of the other molecule and influence its position. Polar molecules have a net attraction between them. Examples of polar molecules include [[hydrogen chloride]] (HCl) and [[chloroform]] (CHCl<sub>3</sub>).
-Also called [[London forces]] or Van der Waals forces, these involve the attraction between temporaily induced dipoles in nonpolar molecules.  This polarization can be induced either by a polar molecule or by the repulsion of negatively charged electron clouds in nonpolar molecules.  An example of the former is chlorine dissolving in water:
+:{{Dipole-dipole-interaction-in-HCl-2D}}
-                  (+)(-)(+)  (-) (+)
- [Permanent Dipole] H-O-H-----Cl-Cl [Induced Dipole]
+Often molecules contain dipolar groups of atoms, but have no overall [[Electric dipole moment|dipole moment]] on the molecule as a whole. This occurs if there is symmetry within the molecule that causes the dipoles to cancel each other out. This occurs in molecules such as [[tetrachloromethane]] and [[carbon dioxide]]. The dipole–dipole interaction between two individual atoms is usually zero, since atoms rarely carry a permanent dipole.
-An example of the second scenario is found in molecular chlorine:
+The Keesom interaction is a van der Waals force. It is discussed further in the section "Van der Waals forces".
-                (+) (-)    (+) (-)
- [Induced Dipole]Cl-Cl------Cl-Cl [Induced Dipole]
+===Ion–dipole and ion–induced dipole forces===
-HI
+Ion–dipole and ion–induced dipole forces are similar to dipole–dipole and dipole–induced dipole interactions but involve ions, instead of only polar and non-polar molecules. Ion–dipole and ion–induced dipole forces are stronger than dipole–dipole interactions because the charge of any ion is much greater than the charge of a dipole moment. Ion–dipole bonding is stronger than hydrogen bonding.<ref>{{Cite book|title=Chemistry: A Molecular Approach| vauthors = Tro N |publisher=Pearson Education Inc|year=2011|isbn=978-0-321-65178-5|location=United States|pages=466}}</ref>
-''See also:'' [[polymer]]
+An ion–dipole force consists of an ion and a polar molecule interacting. They align so that the positive and negative groups are next to one another, allowing maximum attraction. An important example of this interaction is hydration of ions in water which give rise to [[Hydration energy|hydration enthalpy]]. The polar water molecules surround themselves around ions in water and the energy released during the process is known as hydration enthalpy. The interaction has its immense importance in justifying the stability of various ions (like Cu<sup>2+</sup>) in water.
+An ion–induced dipole force consists of an ion and a non-polar molecule interacting. Like a dipole–induced dipole force, the charge of the ion causes distortion of the electron cloud on the non-polar molecule.<ref name=Michael-Blaber-1996>{{cite web | vauthors = Blaber M | date = 1996 | url = http://www.mikeblaber.org/oldwine/chm1045/notes/Forces/Intermol/Forces02.htm | title = Intermolecular Forces | work = mikeblaber.org | access-date = 2011-11-17 | archive-date = 2020-08-01 | archive-url = https://web.archive.org/web/20200801205131/http://www.mikeblaber.org/oldwine/chm1045/notes/Forces/Intermol/Forces02.htm | url-status = usurped }}</ref>
+==Van der Waals forces==
+{{Main|van der Waals force}}
+The van der Waals forces arise from interaction between uncharged atoms or molecules, leading not only to such phenomena as the cohesion of condensed phases and physical absorption of gases, but also to a universal force of attraction between macroscopic bodies.<ref name=j1>{{cite journal | vauthors = Leite FL, Bueno CC, Da Róz AL, Ziemath EC, Oliveira ON | title = Theoretical models for surface forces and adhesion and their measurement using atomic force microscopy | journal = International Journal of Molecular Sciences | volume = 13 | issue = 10 | pages = 12773–12856 | date = October 2012 | pmid = 23202925 | pmc = 3497299 | doi = 10.3390/ijms131012773 | doi-access = free }}</ref>
+===Keesom force (permanent dipole – permanent dipole) {{Anchor|Keesom force}}===
+The first contribution to van der Waals forces is due to electrostatic interactions between rotating permanent dipoles, quadrupoles (all molecules with symmetry lower than cubic), and  multipoles.  It is termed the ''Keesom interaction'', named after [[Willem Hendrik Keesom]].<ref>{{cite journal| vauthors = Keesom WH |title=The second virial coefficient for rigid spherical molecules whose mutual attraction is equivalent to that of a quadruplet placed at its center |url=http://www.dwc.knaw.nl/DL/publications/PU00012540.pdf |journal= Proceedings of the Royal Netherlands Academy of Arts and Sciences |year=1915 |volume= 18 |pages= 636–646}}</ref> These forces originate from the attraction between permanent dipoles (dipolar molecules) and are temperature dependent.<ref name=j1/>
+They consist of attractive interactions between dipoles that are [[canonical ensemble|ensemble]] averaged over different rotational orientations of the dipoles. It is assumed that the molecules are constantly rotating and never get locked into place. This is a good assumption, but at some point molecules do get locked into place. The energy of a Keesom interaction depends on the inverse sixth power of the distance, unlike the interaction energy of two spatially fixed dipoles, which depends on the inverse third power of the distance. The Keesom interaction can only occur among molecules that possess permanent dipole moments, i.e., two polar molecules. Also Keesom interactions are very weak van der Waals interactions and do not occur in aqueous solutions that contain electrolytes. The angle averaged interaction is given by the following equation:
+:<math>\frac{-d_1^2 d_2^2}{24\pi^2 \varepsilon_0^2 \varepsilon_r^2 k_\text{B} T r^6} = V,</math>
+where ''d'' = electric dipole moment, <math>\varepsilon_0</math> = permittivity of free space, <math>\varepsilon_r</math> = dielectric constant of surrounding material, ''T'' = temperature, <math>k_\text{B}</math> = Boltzmann constant, and ''r'' = distance between molecules.
+===Debye force (permanent dipoles–induced dipoles) {{Anchor|Debye force}}===
+The second contribution is the induction (also termed polarization) or Debye force, arising from interactions between rotating permanent dipoles and from the polarizability of atoms and molecules (induced dipoles). These induced dipoles occur when one molecule with a permanent dipole repels another molecule's electrons. A molecule with permanent dipole can induce a dipole in a similar neighboring molecule and cause mutual attraction. Debye forces cannot occur between atoms. The forces between induced and permanent dipoles are not as temperature dependent as Keesom interactions because the induced dipole is free to shift and rotate around the polar molecule. The Debye induction effects and Keesom orientation effects are termed polar interactions.<ref name=j1/>
+The induced dipole forces appear from the induction (also termed [[Dipolar polarization|polarization]]), which is the attractive interaction between a permanent multipole on one molecule with an induced (by the former di/multi-pole) 31 on another.<ref name=Blustin-1978>{{Cite journal |doi= 10.1007/BF00577166 |title= A Floating Gaussian Orbital calculation on argon hydrochloride (Ar·HCl) |journal= Theoretica Chimica Acta |volume= 47 |issue= 3 |pages= 249–257 |year= 1978 | vauthors = Blustin PH |s2cid= 93104668 }}</ref><ref name=Roberts-Orr-1938>{{Cite journal |doi= 10.1039/TF9383401346 |title= Induced dipoles and the heat of adsorption of argon on ionic crystals |journal= Transactions of the Faraday Society |volume= 34 |pages= 1346 |year= 1938 | vauthors = Roberts JK, Orr WJ }}</ref><ref name=Sapse-et-al-1979>{{Cite journal |doi= 10.1038/278332a0 |title= Ion-induced dipole H−n clusters |journal= Nature |volume= 278 |issue= 5702 |pages= 332–333 |year= 1979 | vauthors = Sapse AM, Rayez-Meaume MT, Rayez JC, Massa LJ |bibcode= 1979Natur.278..332S|s2cid= 4304250 }}</ref> This interaction is called the ''Debye force'', named after [[Peter J. W. Debye]].
+One example of an induction interaction between permanent dipole and induced dipole is the interaction between HCl and Ar. In this system, Ar experiences a dipole as its electrons are attracted (to the H side of HCl) or repelled (from the Cl side) by HCl.<ref name=Blustin-1978/><ref name=Roberts-Orr-1938/> The angle averaged interaction is given by the following equation:
+:<math>\frac{-d_1^2 \alpha_2}{16\pi^2 \varepsilon_0^2 \varepsilon_r^2 r^6} = V,</math>
+where <math>\alpha_2</math> = polarizability.
+This kind of interaction can be expected between any polar molecule and non-polar/symmetrical molecule. The induction-interaction force is far weaker than dipole–dipole interaction, but stronger than the [[London dispersion force]].
+===London dispersion force (fluctuating dipole–induced dipole interaction)===
+{{Main|London dispersion force}}
+The third and dominant contribution is the dispersion or London force (fluctuating dipole–induced dipole), which arises due to the non-zero instantaneous dipole moments of all atoms and molecules. Such polarization can be induced either by a polar molecule or by the repulsion of negatively charged electron clouds in non-polar molecules. Thus, London interactions are caused by random fluctuations of electron density in an electron cloud.  An atom with a large number of electrons will have a greater associated London force than an atom with fewer electrons. The dispersion (London) force is the most important component because all materials are polarizable, whereas Keesom and Debye forces require permanent dipoles. The London interaction is universal and is present in atom-atom interactions as well. For various reasons, London interactions (dispersion) have been considered relevant for interactions between macroscopic bodies in condensed systems. [[Hamaker theory|Hamaker]] developed the theory of van der Waals between macroscopic bodies in 1937 and showed that the additivity of these interactions renders them considerably more long-range.<ref name=j1/>
+==Relative strength of forces==
+{|class="wikitable"
+!Bond type
+!Dissociation energy<br />(kcal/mol)<ref name=Seyhan-Organic-Chemistry>{{cite book | vauthors = Eğe SN |title=Organic Chemistry: Structure and Reactivity |date=2004 |publisher=Houghton Mifflin Company |location=Boston |isbn=978-0-618-31809-4 |edition=5th | pages = 30–33, 67}}</ref>
+!Dissociation energy
+(kJ/mol)
+!Note
+|-
+|Ionic lattice
+|250–4000<ref name=Purdue-Lattice>{{cite web |title=Lattice Energies |url=http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch7/lattice.html |access-date=2014-01-21 | work = Division of Chemical Education | publisher = Purdue University }}</ref>
+|1100–20000
+|
+|-
+|Covalent bond
+|30–260
+|130–1100
+|
+|-
+|[[Hydrogen bond]]
+|1–12
+|4–50
+|About 5&nbsp;kcal/mol (21&nbsp;kJ/mol) in water
+|-
+|Dipole–dipole
+|0.5–2
+|2–8
+|
+|-
+|London dispersion forces
+|<1 to 15
+|<4 to 63
+|Estimated from the enthalpies of vaporization of hydrocarbons<ref name=Majer-Svoboda-enthalpy-vap>{{cite book | vauthors = Majer V, Svoboda V |title=Enthalpies of Vaporization of Organic Compounds |date=1985 |publisher=Blackwell Scientific |location=Oxford |isbn=978-0-632-01529-0}}</ref>
+|}
+This comparison is approximate. The actual relative strengths will vary depending on the molecules involved. For instance, the presence of water creates competing interactions that greatly weaken the strength of both ionic and hydrogen bonds.<ref>{{Cite book |last1=Alberts |first1=Bruce |display-authors=etal |url=https://www.worldcat.org/oclc/887605755 |title=Molecular biology of the cell |date=2015 |isbn=978-0-8153-4432-2 |edition=6th |publisher=[[Garland Science]] - [[Taylor & Francis]] |location=New York, NY |oclc=887605755 }}</ref> We may consider that for static systems, [[Ionic bonding]] and [[covalent bond]]ing will always be stronger than intermolecular forces in any given substance. But it is not so for big moving systems like [[enzyme]] molecules interacting with [[Substrate (chemistry)|substrate]] molecules.<ref>{{cite journal | vauthors = Savir Y, Tlusty T | title = Conformational proofreading: the impact of conformational changes on the specificity of molecular recognition | journal = PLOS ONE | volume = 2 | issue = 5 | pages = e468 | date = May 2007 | pmid = 17520027 | pmc = 1868595 | doi = 10.1371/journal.pone.0000468 | bibcode = 2007PLoSO...2..468S | doi-access = free }}</ref> Here the numerous intramolecular (most often - [[hydrogen bond]]s) bonds form an active intermediate state where the intermolecular bonds cause some of the [[covalent bond]] to be broken, while the others are formed, in this way proceeding the thousands of [[Enzyme catalysis|enzymatic reactions]], so important for [[Organism|living organisms]].
+==Effect on the behavior of gases==
+Intermolecular forces are repulsive at short distances and attractive at long distances (see the [[Lennard-Jones potential]]).<ref>{{Cite journal |last1=Fischer |first1=Johann |last2=Wendland |first2=Martin |date=October 2023 |title=On the history of key empirical intermolecular potentials |journal=Fluid Phase Equilibria |language=en |volume=573 |pages=113876 |doi=10.1016/j.fluid.2023.113876|bibcode=2023FlPEq.57313876F |doi-access=free }}</ref><ref>{{Cite journal |last1=Lenhard |first1=Johannes |last2=Stephan |first2=Simon |last3=Hasse |first3=Hans |date=June 2024 |title=On the History of the Lennard-Jones Potential |url=https://onlinelibrary.wiley.com/doi/10.1002/andp.202400115 |journal=Annalen der Physik |language=en |volume=536 |issue=6 |doi=10.1002/andp.202400115 |issn=0003-3804}}</ref> In a gas, the repulsive force chiefly has the effect of keeping two molecules from occupying the same volume. This gives a [[real gas]] a tendency to occupy a larger volume than an [[ideal gas]] at the same temperature and pressure. The attractive force draws molecules closer together and gives a real gas a tendency to occupy a smaller volume than an ideal gas. Which interaction is more important depends on temperature and pressure (see [[compressibility factor]]).
+In a gas, the distances between molecules are generally large, so intermolecular forces have only a small effect. The attractive force is not overcome by the repulsive force, but by the [[thermal energy]] of the molecules. [[Thermodynamic temperature|Temperature]] is the measure of thermal energy, so increasing temperature reduces the influence of the attractive force. In contrast, the influence of the repulsive force is essentially unaffected by temperature.
+When a gas is compressed to increase its density, the influence of the attractive force increases. If the gas is made sufficiently dense, the attractions can become large enough to overcome the tendency of thermal motion to cause the molecules to disperse. Then the gas can condense to form a solid or liquid, i.e., a condensed phase. Lower temperature favors the formation of a condensed phase. In a condensed phase, there is very nearly a balance between the attractive and repulsive forces.
+==Quantum mechanical theories==
+{{Main|Covalent bond#Quantum mechanical description}}
+Intermolecular forces observed between atoms and molecules can be described phenomenologically as occurring between permanent and instantaneous dipoles, as outlined above. Alternatively, one may seek a fundamental, unifying theory that is able to explain the various types of interactions such as [[hydrogen bonding]],<ref name=":0">{{Cite journal| vauthors = Arunan E, Desiraju GR, Klein RA, Sadlej J, Scheiner S, Alkorta I, Clary DC, Crabtree RH, Dannenberg JJ, Hobza P, Kjaergaard HG | display-authors = 6 |date=2011-07-08|title=Definition of the hydrogen bond (IUPAC Recommendations 2011)|journal=Pure and Applied Chemistry|volume=83|issue=8|pages=1637–1641|doi=10.1351/PAC-REC-10-01-02|s2cid=97688573|issn=1365-3075|doi-access=free}}</ref> [[van der Waals force]]<ref name=LD>{{cite book | vauthors = Landau LD, Lifshitz EM | title = Electrodynamics of Continuous Media | url = https://archive.org/details/electrodynamicso00land | url-access = registration | publisher = Pergamon | location = Oxford | date = 1960 | pages = [https://archive.org/details/electrodynamicso00land/page/368 368–376] }}</ref> and dipole–dipole interactions. Typically, this is done by applying the ideas of [[quantum mechanics]] to molecules, and Rayleigh–Schrödinger [[perturbation theory]] has been especially effective in this regard. When applied to existing [[quantum chemistry]] methods, such a quantum mechanical explanation of intermolecular interactions provides an array of approximate methods that can be used to analyze intermolecular interactions.<ref>{{cite journal | doi = 10.1021/ja00428a004 | title = Theory of the Chemical Bond | year = 1976 | vauthors = King M | journal = Journal of the American Chemical Society | volume = 98 | issue = 12 | pages = 3415–3420 }}</ref>  One of the most helpful methods to visualize this kind of intermolecular interactions, that we can find in quantum chemistry, is the [[Non-covalent interactions index|non-covalent interaction index]], which is based on the electron density of the system. London dispersion forces play a big role with this.
+Concerning electron density topology, recent methods based on electron density gradient methods have emerged recently, notably with the development of IBSI (Intrinsic Bond Strength Index),<ref>{{cite journal | vauthors = Klein J, Khartabil H, Boisson JC, Contreras-García J, Piquemal JP, Hénon E | title = New Way for Probing Bond Strength | journal = The Journal of Physical Chemistry A | volume = 124 | issue = 9 | pages = 1850–1860 | date = March 2020 | pmid = 32039597 | doi = 10.1021/acs.jpca.9b09845 | s2cid = 211070812 | bibcode = 2020JPCA..124.1850K | url = https://hal.archives-ouvertes.fr/hal-02377737/file/4_IGM_bond13012020_11h49%20%281%29.pdf }}</ref> relying on the IGM (Independent Gradient Model) methodology.<ref>{{cite journal | vauthors = Lefebvre C, Rubez G, Khartabil H, Boisson JC, Contreras-García J, Hénon E | title = Accurately extracting the signature of intermolecular interactions present in the NCI plot of the reduced density gradient versus electron density | journal = Physical Chemistry Chemical Physics | volume = 19 | issue = 27 | pages = 17928–17936 | date = July 2017 | pmid = 28664951 | doi = 10.1039/C7CP02110K | bibcode = 2017PCCP...1917928L | url = https://hal.univ-reims.fr/hal-02505160/file/2017_IGM_PROMOL%20%281%29.pdf }}</ref><ref>{{cite journal | vauthors = Lefebvre C, Khartabil H, Boisson JC, Contreras-García J, Piquemal JP, Hénon E | title = The Independent Gradient Model: A New Approach for Probing Strong and Weak Interactions in Molecules from Wave Function Calculations | journal = ChemPhysChem | volume = 19 | issue = 6 | pages = 724–735 | date = March 2018 | pmid = 29250908 | doi = 10.1002/cphc.201701325 | url = https://hal.univ-reims.fr/hal-03377532/file/37_version.pdf }}</ref><ref>{{cite journal | vauthors = Ponce-Vargas M, Lefebvre C, Boisson JC, Hénon E | title = Atomic Decomposition Scheme of Noncovalent Interactions Applied to Host-Guest Assemblies | journal = Journal of Chemical Information and Modeling | volume = 60 | issue = 1 | pages = 268–278 | date = January 2020 | pmid = 31877034 | doi = 10.1021/acs.jcim.9b01016 | s2cid = 209488458 }}</ref>
+== See also ==
+{{div col|colwidth=18em}}
+* [[Ionic bonding]]
+* [[Salt bridge (protein and supramolecular)|Salt bridges]]
+* [[Coomber's relationship]]
+* [[Force field (chemistry)]]
+* [[Hydrophobic effect]]
+* [[Intramolecular force]]
+* [[Molecular solid]]
+* [[Polymer]]
+* [[Quantum chemistry computer programs]]
+* [[van der Waals force]]
+* [[Comparison of software for molecular mechanics modeling]]
+* [[Non-covalent interactions]]
+* [[Solvation]]
+{{div col end}}
+== References ==
+{{Reflist|30em}}
+{{Chemical bonds}}{{Authority control}}
+{{DEFAULTSORT:Intermolecular Force}}
+[[Category:Intermolecular forces]]
+[[Category:Chemical bonding]]
+[[Category:Johannes Diderik van der Waals]]