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An '''intermolecular force''' ('''IMF''') (or '''secondary force''') is the force that mediates interaction between molecules, including the [[Electromagnetism|electromagnetic forces of attraction
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 |url=https://linkinghub.elsevier.com/retrieve/pii/S0378381223001565 |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]]
Attractive intermolecular forces are categorized into the following types:
*[[Hydrogen bond]]ing
*Ion–dipole forces and ion–induced dipole
*[[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]]
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]]).
==Hydrogen bonding==▼
▲== Hydrogen bonding ==
{{Main|Hydrogen bond}}
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[[Image:Hydrogen-bonding-in-water-2D.svg|thumb|Hydrogen bonding in water]]
Though both not depicted in the diagram, water molecules have
==
{{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>
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.<ref>{{cite journal | vauthors = Biedermann F, Schneider HJ | title = Experimental
==Dipole–dipole and similar interactions {{anchor|Dipole-dipole interactions}}==
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:{{Dipole-dipole-interaction-in-HCl-2D}}
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.
The Keesom interaction is a van der Waals force. It is discussed further in the section "Van der Waals forces".
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===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
:<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> =
===Debye force (permanent dipoles–induced dipoles) {{Anchor|Debye force}}===
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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 |
==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 |url=https://linkinghub.elsevier.com/retrieve/pii/S0378381223001565 |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.
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{{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|
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>{{
== See also ==
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