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'''Intermolecular forces''' are forces of attraction or repulsion which act between neighboring particles (atoms, molecules or ions). They are weak compared to the [[intramolecular force]]s, the forces which keep a molecule together. For example, the [[covalent bond]] present within HCl molecules is much stronger than the forces present between the neighboring molecules, which exist when the molecules are sufficiently close to each other. | |||
If a bond has a H atom directly bonded to either O, N, or F- then it is an H-Bond. | |||
If it is polar, then it is a dipole; if it is nonpolar, then it is a dispersion force (London). | |||
The investigation of intermolecular forces starts from macroscopic observations which point out the existence and action of forces at molecular level. These observations include non-ideal gas thermodynamic behavior reflected by [[virial coefficient]]s, [[vapor pressure]], [[viscosity]], superficial tension and adsorption data. | |||
The first reference to the nature of microscopic forces is found in [[Alexis Clairaut]]'s work Theorie de la Figure de la Terre.<ref>H. Margenau, N Kestner, Theory of intermolecular forces, International Series of Monographs in Natural Philosophy, Pergamon Press</ref> Other scientists who have contributed to the investigation of microscopic forces include: [[Laplace]], [[Gauss]], [[J. C. Maxwell|Maxwell]] and [[Boltzmann]]. | |||
Attractive intermolecular forces are considered by the following types: | |||
*Dipole-dipole forces | |||
*Ion-dipole forces | |||
*[[Van der Waals force]]s ([[Van der Waals forces#Keesom force|Keesom force]], [[Debye force]], and [[London dispersion force]]) | |||
Information on intermolecular force is obtained by macroscopic measurements of properties like viscosity, PVT data. The link to microscopic aspects is given by [[virial coefficient]]s and [[Lennard-Jones potential]]s. | |||
==Dipole-dipole interactions {{anchor|Dipole-dipole interactions}} == | |||
Dipole-dipole interactions are electrostatic interactions between permanent [[dipole]]s in molecules. 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>). | |||
[[Image:Dipole-dipole-interaction-in-HCl-2D.png|200px]] | |||
Often molecules contain dipolar groups, but have no overall [[Molecular dipole moment|dipole moment]]. 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]]. Note that the dipole-dipole interaction between two individual atoms is usually zero, since atoms rarely carry a permanent dipole. See [[Dipole#Atomic dipoles|atomic dipoles]]. | |||
=== Ion-dipole and ion-induced dipole forces === | |||
Ion-dipole and ion-induced dipole forces are similar to dipole-dipole and 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.{{citation needed|date=July 2012}} | |||
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 for maximum attraction. | |||
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>Dr. Michael Blaber, 1996. Intermolecular Forces. http://www.mikeblaber.org/oldwine/chm1045/notes/Forces/Intermol/Forces02.htm</ref> | |||
===Hydrogen bonding=== | |||
{{Main|Hydrogen bond}} | |||
A '''hydrogen bond''' is the attraction between the lone pair of an [[electronegative]] atom and a [[hydrogen]] atom that is bonded to either [[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|van der Waals interaction]], produces interatomic distances shorter than the sum of [[van der Waals radius]], and usually involves a limited number of interaction partners, which can be interpreted as a kind of [[Valence (chemistry)|valence]]. | |||
[[Image:Hydrogen-bonding-in-water-2D.png|200px]] | |||
Intermolecular hydrogen bonding is responsible for the high boiling point of [[water]] (100 °C) compared to the other [[chalcogen|group 16]] [[hydride]]s, which have no hydrogen bonds. 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.{{citation needed|date=July 2012}} | |||
==Van der Waals forces== | |||
{{Main|Van der Waals force}} | |||
===Keesom (permanent dipole) force=== | |||
'''Keesom interactions''' (named after [[Willem Hendrik Keesom]]) are attractive interactions of 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 aka 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{-2m_1^2m_2^2}{48\pi^2\varepsilon_o^2\varepsilon_r^2k_bTr^6}=V</math> | |||
Where m = charge per length, <math>\varepsilon_o</math> = permitivity of free space, <math>\varepsilon_r</math> = dielectric constant of surrounding material, T = temperature, <math>k_b</math> = Boltzmann constant, and r = distance between molecules. | |||
===Debye (induced dipole) force=== | |||
The induced dipole forces appear from the induction (also known as [[Dipolar polarization|polarization]]), which is the attractive interaction between a permanent multipole on one molecule with an induced (by the former di/multi-pole) multipole on another.<ref name=Blustin-1978>Blustin PH, 1978. A Floating Gaussian Orbital calculation on argon hydrochloride (Ar • HCl). Theoret. Chim. Acta 47, 249–257.</ref><ref name=Nannoolal-2006>Nannoolal Y, 2006. Development and critical evaluation of group contribution methods for the estimation of critical properties, liquid vapour pressure and liquid viscosity of organic compounds. University of Kwazulu-Natal PhD Thesis.</ref><ref name=Roberts-Orr-1938>Roberts JK and Orr WJC, 1938. Induced dipoles and the heat of adsorption of argon on ionic crystals. Trans. Faraday Soc. 34, 1346–1349.</ref><ref name=Sapse-et-al-1979>Sapse AM, Rayez-Meaume MT, Rayez JC and Massa LJ, 1979. Ion-induced dipole H-n clusters. Nature 278, 332–333.</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{-m_1^2\alpha_2}{16\pi^2\varepsilon_o^2\varepsilon_r^2r^6}=V</math> | |||
Where <math>\alpha</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 === | |||
{{Main|London dispersion force}} | |||
Otherwise known as quantum-induced instantaneous polarization or instantaneous dipole-induced dipole force, the [[London dispersion force]] is caused by [[Electronic correlation|correlated movements of the electrons]] in interacting molecules. Electrons that belong to different molecules start "fleeing" and avoiding each other at the short intermolecular distances, which is frequently described as formation of "instantaneous dipoles" that attract each other. | |||
==Relative strength of forces== | |||
{|class="wikitable" | |||
!Bond type | |||
!Dissociation energy (kcal/mol)<ref name=Seyhan-Organic-Chemistry>Organic Chemistry: Structure and Reactivity by Seyhan Ege, pp.30–33, 67</ref> | |||
|- | |||
|Ionic Lattice Energy | |||
|250-4000 <ref name=Purdue-Lattice>{{cite web |title=Lattice Energies |url=http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch7/lattice.html |accessdate=2014-01-21}}</ref> | |||
|- | |||
|Covalent Bond Energy | |||
|30-260 | |||
|- | |||
|London Dispersion Forces | |||
|<1 to 15 (estimated from the enthalpies of vaporization of hydrocarbons)<ref name=Majer-Svoboda-enthalpy-vap>Majer, V. and Svoboda, V., Enthalpies of Vaporization of Organic Compounds, Blackwell Scientific Publications, Oxford, 1985.}</ref> | |||
|- | |||
|[[Hydrogen bond|Hydrogen Bonds]] | |||
|1-12 (about 5 in water) | |||
|- | |||
|Dipole–Dipole | |||
|0.5–2 {{citation needed|date=January 2014}} | |||
|} | |||
Note: this comparison is only approximate – the actual relative strengths will vary depending on the molecules involved. Ionic and covalent bonding will always be stronger than intermolecular forces in any given substance. For very small, highly polar molecules with hydrogen bonding, London Dispersion forces may be weaker than hydrogen bonds; however for most molecules, the [[London dispersion force]] will be the dominant intermolecular force affecting their properties (even for ammonia, this is estimated to account for more than 50% of the total attractive force between molecules). | |||
==Quantum mechanical theories== | |||
{{Main|Quantum mechanical explanation of intermolecular interactions}} | |||
{{Expand section|date=September 2009}} | |||
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, van der Waals forces 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]], this provides an array of approximate methods that can be used to analyze intermolecular interactions. | |||
== See also == | |||
{{Div col|2}} | |||
* [[Coomber's relationship]] | |||
* [[Force field (chemistry)|Force field]] | |||
* [[Hydrophobic effect]] | |||
* [[Intramolecular force]] | |||
* [[Molecular solid]] | |||
* [[Polymer]] | |||
* [[Quantum chemistry computer programs]] | |||
* [[Van der Waals force]] | |||
* [[List of software for molecular mechanics modeling|Software for molecular mechanics modeling]] | |||
* [[Non-covalent interactions]] | |||
* [[Solvation]] | |||
{{Div col end}} | |||
==References== | |||
<references/> | |||
==External links== | |||
;Software for calculation of intermolecular forces | |||
*[http://www.q-pharm.com Quantum 3.2] | |||
*[http://www.physics.udel.edu/~szalewic/SAPT/ SAPT]: An ab initio quantumchemical package | |||
*[http://scitation.aip.org/content/aip/journal/jcp/105/20/10.1063/1.472747 intermolecular potential] | |||
{{Chemical bonds}} | |||
{{DEFAULTSORT:Intermolecular Force}} | |||
[[Category:Intermolecular forces]] | |||
[[Category:Chemical bonding]] | |||
[[ar:ترابط بين جزئي خارجي]] | |||
[[zh:分子间作用力]] | |||
Revision as of 12:29, 14 January 2014
Intermolecular forces are forces of attraction or repulsion which act between neighboring particles (atoms, molecules or ions). They are weak compared to the intramolecular forces, the forces which keep a molecule together. For example, the covalent bond present within HCl molecules is much stronger than the forces present between the neighboring molecules, which exist when the molecules are sufficiently close to each other.
If a bond has a H atom directly bonded to either O, N, or F- then it is an H-Bond. If it is polar, then it is a dipole; if it is nonpolar, then it is a dispersion force (London).
The investigation of intermolecular forces starts from macroscopic observations which point out the existence and action of forces at molecular level. These observations include non-ideal gas thermodynamic behavior reflected by virial coefficients, vapor pressure, viscosity, superficial tension and adsorption data.
The first reference to the nature of microscopic forces is found in Alexis Clairaut's work Theorie de la Figure de la Terre.[1] Other scientists who have contributed to the investigation of microscopic forces include: Laplace, Gauss, Maxwell and Boltzmann.
Attractive intermolecular forces are considered by the following types:
- Dipole-dipole forces
- Ion-dipole forces
- Van der Waals forces (Keesom force, Debye force, and London dispersion force)
Information on intermolecular force is obtained by macroscopic measurements of properties like viscosity, PVT data. The link to microscopic aspects is given by virial coefficients and Lennard-Jones potentials.
Dipole-dipole interactions <Dipole-dipole interactions>...</Dipole-dipole interactions>
Dipole-dipole interactions are electrostatic interactions between permanent dipoles in molecules. 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).
Often molecules contain dipolar groups, but have no overall dipole moment. 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. Note that the dipole-dipole interaction between two individual atoms is usually zero, since atoms rarely carry a permanent dipole. See atomic dipoles.
Ion-dipole and ion-induced dipole forces
Ion-dipole and ion-induced dipole forces are similar to dipole-dipole and 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.Potter or Ceramic Artist Truman Bedell from Rexton, has interests which include ceramics, best property developers in singapore developers in singapore and scrabble. Was especially enthused after visiting Alejandro de Humboldt National Park.
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 for maximum attraction.
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.[2]
Hydrogen bonding
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A hydrogen bond is the attraction between the lone pair of an electronegative atom and a hydrogen atom that is bonded to either nitrogen, oxygen, or fluorine.[3] 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 interaction, produces interatomic distances shorter than the sum of van der Waals radius, and usually involves a limited number of interaction partners, which can be interpreted as a kind of valence.
Intermolecular hydrogen bonding is responsible for the high boiling point of water (100 °C) compared to the other group 16 hydrides, which have no hydrogen bonds. 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.Potter or Ceramic Artist Truman Bedell from Rexton, has interests which include ceramics, best property developers in singapore developers in singapore and scrabble. Was especially enthused after visiting Alejandro de Humboldt National Park.
Van der Waals forces
Mining Engineer (Excluding Oil ) Truman from Alma, loves to spend time knotting, largest property developers in singapore developers in singapore and stamp collecting. Recently had a family visit to Urnes Stave Church.
Keesom (permanent dipole) force
Keesom interactions (named after Willem Hendrik Keesom) are attractive interactions of 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 aka 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:
Where m = charge per length, = permitivity of free space, = dielectric constant of surrounding material, T = temperature, = Boltzmann constant, and r = distance between molecules.
Debye (induced dipole) force
The induced dipole forces appear from the induction (also known as polarization), which is the attractive interaction between a permanent multipole on one molecule with an induced (by the former di/multi-pole) multipole on another.[4][5][6][7] 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.[4][6] The angle averaged interaction is given by the following equation.
Where = 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
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Otherwise known as quantum-induced instantaneous polarization or instantaneous dipole-induced dipole force, the London dispersion force is caused by correlated movements of the electrons in interacting molecules. Electrons that belong to different molecules start "fleeing" and avoiding each other at the short intermolecular distances, which is frequently described as formation of "instantaneous dipoles" that attract each other.
Relative strength of forces
| Bond type | Dissociation energy (kcal/mol)[8] |
|---|---|
| Ionic Lattice Energy | 250-4000 [9] |
| Covalent Bond Energy | 30-260 |
| London Dispersion Forces | <1 to 15 (estimated from the enthalpies of vaporization of hydrocarbons)[10] |
| Hydrogen Bonds | 1-12 (about 5 in water) |
| Dipole–Dipole | 0.5–2 Potter or Ceramic Artist Truman Bedell from Rexton, has interests which include ceramics, best property developers in singapore developers in singapore and scrabble. Was especially enthused after visiting Alejandro de Humboldt National Park. |
Note: this comparison is only approximate – the actual relative strengths will vary depending on the molecules involved. Ionic and covalent bonding will always be stronger than intermolecular forces in any given substance. For very small, highly polar molecules with hydrogen bonding, London Dispersion forces may be weaker than hydrogen bonds; however for most molecules, the London dispersion force will be the dominant intermolecular force affecting their properties (even for ammonia, this is estimated to account for more than 50% of the total attractive force between molecules).
Quantum mechanical theories
Mining Engineer (Excluding Oil ) Truman from Alma, loves to spend time knotting, largest property developers in singapore developers in singapore and stamp collecting. Recently had a family visit to Urnes Stave Church.
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, van der Waals forces 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, this provides an array of approximate methods that can be used to analyze intermolecular interactions.
See also
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- Coomber's relationship
- Force field
- Hydrophobic effect
- Intramolecular force
- Molecular solid
- Polymer
- Quantum chemistry computer programs
- Van der Waals force
- Software for molecular mechanics modeling
- Non-covalent interactions
- Solvation
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References
- ↑ H. Margenau, N Kestner, Theory of intermolecular forces, International Series of Monographs in Natural Philosophy, Pergamon Press
- ↑ Dr. Michael Blaber, 1996. Intermolecular Forces. http://www.mikeblaber.org/oldwine/chm1045/notes/Forces/Intermol/Forces02.htm
- ↑ Template:GoldBookRef
- ↑ 4.0 4.1 Blustin PH, 1978. A Floating Gaussian Orbital calculation on argon hydrochloride (Ar • HCl). Theoret. Chim. Acta 47, 249–257.
- ↑ Nannoolal Y, 2006. Development and critical evaluation of group contribution methods for the estimation of critical properties, liquid vapour pressure and liquid viscosity of organic compounds. University of Kwazulu-Natal PhD Thesis.
- ↑ 6.0 6.1 Roberts JK and Orr WJC, 1938. Induced dipoles and the heat of adsorption of argon on ionic crystals. Trans. Faraday Soc. 34, 1346–1349.
- ↑ Sapse AM, Rayez-Meaume MT, Rayez JC and Massa LJ, 1979. Ion-induced dipole H-n clusters. Nature 278, 332–333.
- ↑ Organic Chemistry: Structure and Reactivity by Seyhan Ege, pp.30–33, 67
- ↑ Template:Cite web
- ↑ Majer, V. and Svoboda, V., Enthalpies of Vaporization of Organic Compounds, Blackwell Scientific Publications, Oxford, 1985.}
External links
- Software for calculation of intermolecular forces
- Quantum 3.2
- SAPT: An ab initio quantumchemical package
- intermolecular potential