Talk:Intermolecular force

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Sodium ion/water ion-dipole bounding energy[edit]

I found on an article written by teachers saying: "The sodium ion/water cluster interaction is approximately 50 KJ/mol." We can use this amount of energy in the interraction to compare with hydrogen bounding's strength and say that ion-dipole are generally stronger than HB. (would fix the citation needed in ion-dipole text) --JuRien (talk) 03:22, 13 May 2016 (UTC)

Too complex and gratuitous linking from other articles[edit]

Seems like someone (successfully) tried to eradicate the instantaneous dipole explanation of the London dispersion force from Wikipedia. I remember there once was a nice Wikipedia page on this subject (or at the very least a link to a nice external page), but now, wherever I click on a link that says "Induced-dipole attraction", "instantaneous dipole", "dispersion"... I get directed to the "Quantum mechanical theory of dispersion forces" section of this ragtag Wikipedia article (see for instance Dipole#Molecular_dipoles, Van_der_Waals_force#London_dispersion_force,...).

Yes, the quantum mechanical second-order perturbation theory is a more rigorous explanation of the dispersion interaction. No, this doesn't render the instantaneous dipole explanation invalid (as the page suggests). Even modern quantum chemists often fall back to the instantaneous dipole explanation because it's a simpler frame of reference to think in. If you don't believe me, read this paper, written by none other than Schroedinger Medal laureate Axel Becke: J. Chem. Phys. 122, 154104 (2005). Note how he derives his theory using the instantaneous dipole moment of the Fermi hole. Then, he refers to the appendix for the formal derivation using second-order perturbation theory. If even the readership of J. Chem. Phys. finds it easier to think in terms of instantaneous dipoles, do you really expect non-experts that read Wikipedia to understand your second-order perturbation theory?

Same with the whole "electrostatic interactions" section. It may all be correct, but it is definitely not the kind of information your average Wikipedian is looking for. Wikipedia is an encyclopedia, not a physical chemistry textbook!!!

On a more constructive note, maybe we could rewrite the current "intermolecular force" page so that it gives only the classical explanations. For each type of intermolecular force, after giving the classical explanation, we'd include a link to a new page "quantum mechanical explanation of intermolecular interactions", that would contain most of the content from the current page (albeit organized a bit better). Of course, this would be a major overhaul, so there probably needs to be some discussion on the idea before we can start implementing it.

OneAhead (talk) 01:08, 22 August 2009 (UTC)

Van der Waal's information incorrect?[edit]

  • Van der Waal's information here has a contrast to what I learnt at school.

What I learnt was Van der Waal's forces occur because of assymetrical distribution of electron clouds due to movement of electrons, instead of polarity. jynx 15:59, 24 November 2005 (UTC)

See dipole, there's a short conventional division of weak intermolecular forces. Debye, Keesom, Van der Waals, etc. all studied these phenomenon from slightly different angles prior quantum mechanics, i guess hence the confuse. One could say that the uncertainty principle is working here, as the dipole moments at any given time should be exactly derivable but aren't so...  ;-)

Oh, no. Attractive Van der Waal's forces are of three general types: dipole-dipole, dipole-induced dipole, and interactions of "instantaneous dipoles" (London's dispersion attractions), according to good textbook definitions.Biophys 22:56, 17 October 2007 (UTC)
yeah ... Biophys is right . GreekAlexander (talk) 15:45, 21 August 2009 (UTC)

HCl as an example for a dipole-dipole reaction[edit]

Might it be a good idea to use a different polar molecule for dipole-dipole interaction than HCl? Namely something without hydrogen. maybe Acetone or CO or any other aprotic polar molecule? Fett0001 02:13, 1 November 2005 (UTC)

Van der Waal's Volume?[edit]

I came here from "van der Waals volume" on the proline page. I found van der Waals, but what's van der Waals volume? The preceding unsigned comment was added by PierreAbbat (talk • contribs) .

I believe van der Waals forces and "van der Waals volume" relate back to the van der Waals equation, which compensates for intermolecular forces and volume. I believe van der Waals volume would be the volume of a gas molecule that is accounted for by the equation. Don't hold me to that. The preceding unsigned comment was added by Mauvila (talk • contribs) .

Two atomic dipoles weakly attract by Van der Waals interaction, bringing the two nuclei closer. When they draw closer together, their electron clouds begin to repel each other. When van der Waals attraction balances this repulsive force, the distance is the Van der Waals radius, and atoms are at their van der Waals volume; in the space-filling molecular models the atoms are usually shown at their van der Waals radii/volumes.

In my opn some of the problems of this argument is that the title "intermolecular interaction" could better be "non covalent interaction" because many interaction are between atoms; van der waals volumes are volumes for atoms inside molecules. 13:26, 28 March 2007 (UTC)roberto90967


This page would really benefit from some real diagrams. - Omegatron 19:59, May 26, 2004 (UTC)

Van der Waal's forces and London Dispersion Forces as a synonym[edit]

I don't think the Van der Waals forces should be used as a synonym for London dispersion forces. In its true definition, Van der Waals forces are those which contribute to the non-volume part of the Van der Waals equation. The preceding unsigned comment was added by (talk • contribs) .

Correct. See my comment above.Biophys 22:58, 17 October 2007 (UTC)

Under the headline of London dispersion forces I changed: "sometimes called the van der Waals force" to "one of the three types of van der Waals forces".

This after I read comments in the "discussion" banner that agreed with me on that London dispersion forces are not synonomous with van der Waals forces.

There's also a discussion on this from [1]

Thγmφ (talk) 21:17, 8 October 2009 (UTC)

Polar molecules not in Symmetrical molecules[edit]

I've been told that in, say CCl4, even though the electro negativities of the Cls cause each C Cl bond to be a dipole, because of it's symmettry, it causes the overall charge to disappear, meaning that the overall molecule isn't polar whatsoever. Now, why? Surely any charge on the outside of a molecule can't cancel a like force also on the outside of a molecule. The preceding unsigned comment was added by (talk • contribs) .

Don't trust me on this one, but my view is that the symmetry of the dipole bonds makes each bond equally susceptible to outside influences, so would there be an environment that effects to these bonds, the molecule still would rotate, unless it'd be 'anchored' from three points, by an outside &delta+, &delta- , and some London forces. This still wouldn't effect the 4th Cl, since partial charges tend cancel themselves out, because on atomic scale they are intentionally divided parts of an Eigenstate of the molecule in question... see also entropy... Had me thinking for a while. I guess quantum chemists could say something more by some formulas... A good question, hope this helps (somethings can be taken as stated).

symmetric molecules do not have a dipole moment, so they are not polar --Spoon! 19:07, 31 August 2006 (UTC)
Not necessarily. Water molecule has a C2 symmetry axis, and it has a dipole moment.Biophys 23:00, 17 October 2007 (UTC)


I reverted this:

Intermolecular forces, the component of the intermolecular bond, are electromagnetic forces...

because the grammar is unsettling and more importantly it adds no useful information. —Keenan Pepper 14:48, 31 January 2006 (UTC)

Merging Keesom force to Intermolecular Force[edit]

I've proposed a merge of the content from Keesom force to Intermolecular Force, because I believe that the term itself is far too unique to be used commonly. The term "Dipole-dipole interaction" has always been used in every textbook I have read, and the term itself redirects here. Keesom force can be made into a redirect. Kareeser|Talk! 00:56, 13 February 2006 (UTC)

More merging may be done with [intermolecular attraction] perhaps? -Kristan Wifler

So there are 5 forces listed:

Attractive intermolecular forces:

   Dipole-dipole forces
   Ion-dipole forces
   Van der Waals forces (Keesom force, Debye force, and London dispersion force)

So Keesom force and dipole-dipole forces are not the same. A simple explanation (catchwords: Boltzmann-averaged, rotating, not fixed) why they differ and why those assumptions are needed/useful, why we need both dipole-dipole interactions respectively would be great. Melierax eng (talk) 19:31, 28 November 2013 (UTC)

clean up[edit]

You bet this article needs a clean up. It coudn't be any more cluttered. I could have made a better article (no I couldn't). Tourskin.

Aren't London dispersion forces, dipole-induced dipole forces, Debye forces, Instantaneous dipole-induced dipole forces, and just plain old instantaneous dipole forces all the same thing? This article needs to be more streamlined by having these combined under the same heading if there are the same. Also, ion-dipole forces are not elaborated on, and this needs to be done so. Make the article more systematic that way. Thanks. (talk) 06:55, 2 December 2010 (UTC)

Possible mistake[edit]

I read in the Hydrogen Bonding section of this article that it occurs when a hydrogen atom is **noncovalently bonded** to an electronegative atom... The fact that it is noncovalently bonded puzzles me. E.g. H is covalently bonded to O in a water molecule. Please explain. --Freiddie 20:49, 5 August 2007 (UTC)

Hydrogen bonding involves a hydrogen atom covalently bonded to one electronegative atom and noncovalently bonded to another electronegative atom. For example, a hydrogen bond between two water molecules can be described diagrammatically as follows: O-H···O. The solid line is a covalent O-H bond and the dashed line is a noncovalent O-H bond. The noncovalent bond is called the hydrogen bond.
Ben 20:54, 5 August 2007 (UTC)
Apparently, I understood this point. But the fact that it is followed by "The result is a dipolar molecule." seems to indicate that the noncovalent bond between this hydrogen and the electronegative atom forms a dipolar molecule. The way the passage describes the hydrogen is really odd here. --Freiddie 18:48, 6 August 2007 (UTC)

Question: is the force attraction of solid,liquid,gas,plasma strong or weak or what?? that is our assignment so please answer —Preceding unsigned comment added by (talk) 10:52, 3 September 2007 (UTC)

This article is terribly useless.[edit]

I think I fell asleep while attempting to read it. People aren't coming to an article about intermolecular forces to see only deep physical theory and accompanying formulae while something as simple as a table comparing types of intermolecular forces is missing. Hell, there's only one mention of kj/mol, and even the most basic chemistry 101 book will tell you the relative strengths of different kinds of intermolecular forces. If an article is going to be chock full of higher-level physics, then it damn well better have the basic, encyclopedic information nailed down.-- 21:28, 19 September 2007 (UTC)

Of course it is not useless, but it definitely needs improvement. Your point is well founded. It should be understandable for a high school student. Perhaps we need to move some parts with a lot of math to smaller and more specialized articles to improve readability. At the same time, some missing qualitative concepts (such as dependence of vdW forces on the environment) should be included. Biophys 23:09, 17 October 2007 (UTC) Yes, Quantum Mechanics treatment of intermolecular forces should be made a separate article. Wikipedia suppose to be for general public. BTW, all intermolecular interactions can be described in framework of classical physics. Even dispersion attractions can be described as interactions of fluctuating dipoles, although they are completely of QM nature (electron-electron correlation).Biophys 05:06, 18 October 2007 (UTC)

Overall critic[edit]

This article is bad from start.

In the first part:

One should start to talk about ideal gases and why the ideal gas model fails to predict intermolecular bonding. This assumption is not present in the text, and is the basis for the entire description of intermolecular forces, either classical, or quantum mechanics.

In physics, chemistry, and biology (Why not just 'In Nature' ?), intermolecular forces are forces that act between stable molecules or between functional groups of macromolecules (It is not specified the kind of 'action', hence, this phrase does not in any away contribute do define anything, and *non-stable* molecules, such as radicals, also show this kind of interactions!). Intermolecular forces (aka van der Waal's forces) include momentary attractions between molecules, diatomic free elements, and individual atoms (There is a need to specify what other kinds of 'interactions' occurs within intermolecular forces, or rewrite the whole phrase). They differ from covalent and ionic bonding in that they are not stable (This phrase must be rewritten. Intermolecular interaction is not so *strong as*, and due to its nature, it's shortlived), but are caused by momentary polarization of particles (Nonsense! Is covalent bond *due to* momentary polarization of particles? Rewrite!). Because electrons have no fixed position in the structure of an atom or molecule, but rather are distributed in a probabilistic fashion based on quantum probability, there is a positive chance that the electrons are not evenly distributed and thus their electrical charges are not evenly distributed. See Schrödinger equation for the theories on wave functions and descriptions of position and velocity of quantum particles.(Rewrite this! This should be shorter and not so extended! If you wish to speak of this, place it in a new subchapter! And by the way, where are the references for van der waals equation? and for the lack of interpretation provided by ideal gas equation? those are a whole lot more important here!)

In general one distinguishes short and long range van der Waal's forces. The former are due to intermolecular exchange and charge penetration (What is intermolecular exchange and charge penetration? One can't just put names! at least put a link to an article that explains it! intermolecular exchange is not so trivial as this part of the text might lead to think!). They fall off exponentially as a function of intermolecular distance R and are repulsive for interacting closed-shell systems (An equation, followed by this explanation would be a nice coming!). In chemistry they are well known(Suggestion: This effects are well-known in physical chemistry, due to the fact that they give rise...), because they give rise to steric hindrance, also known as Born or Pauli repulsion. Long range forces fall off with inverse powers of the distance, R-n, typically 3 ≤ n ≤ 10, and are mostly attractive.(This paragraph needs to be rewritten! It's too confusing!)

The sum of long and short range forces gives rise to a minimum, referred to as Van der Waal minimum (Some equations would be nice! And the references? I haven't seen one single reference YET!). The position and depth of the Van der Waal's minimum depends on distance and mutual orientation of the molecules.(I suggest some figures here, to enhance understanding) "General theory" This is because before the advent of quantum mechanics the origin of intermolecular forces was not well understood. Especially the causes of hard sphere repulsion, postulated by Van der Waals (OH! Only know you speak of the most important part? And what is the hard sphere repulsion? Just a name?), and the possibility of the liquefaction of noble gases were difficult to understand. Soon after the formulation of quantum mechanics, however, all open questions regarding intermolecular forces were answered, first by S.C. Wang and then more completely and thoroughly by Fritz London. (There is no logical historical sequence in this introduction! I strongly recommend a cleanup in this part! Author should mention the ideal gas model as not being able to explain intermolecular forces (such as liquifying gases), van der Waals theory and equation, Keesom work on permanent dipole-permanent-dipole, Debye's work on the permanent-dipole-induced-dipole, and London's QM work on induced dipole, induced dipole. There is a great mixture of concepts and even some concepts that are not even explained! And please, never *ever* try to explain a concept, aplying the same concept, such as 'a force is a force that...)'

More suggestions will come. Right now I am short of time. —Preceding unsigned comment added by Alsimao (talkcontribs) 21:09, 25 May 2008 (UTC)

How about a table comparing the relative strengths of each intermolecular force? That's what I was looking for, but it's nowhere to be found in this article. Perhaps it's hard to make a generalization, as indicated by the disclaimer in the chemical bond template, but this information has much broader applications than the esoteric technical details that currently comprise the bulk of the article.Fuzzform (talk) 06:33, 21 January 2009 (UTC)

Starting the clean up![edit]

Just did a little bit of spring cleaning and tidying up of this article. I added some bits at the start explaining the three models- why this was placed somewhere in the middle of the article I do not know. Because of this I had to comment some of the heavy theory, simply because I don't know what it means so I don't want to mess about with it.

Is anyone else thinking that maybe the maths should be put on another page? I'm thinking students and non-experts (myself included here) will get quite scared by all the equations thrown in their face, without a simple explanation of what an intermolecular force is. Thecurran91 (talk) 19:30, 21 July 2009 (UTC)

I couldn't agree more. As I lamented above, while the classical explanation of intermolecular interactions is not 100% accurate, it is elegant, easy to understand, and leads to conclusions that are correct in a vast majority of cases. The quantum mechanical explanation requires a university-level background in mathematics and quantum physics, and has very little added value from the point of view of a non-theoretical chemist (and I'm a theoretical chemist...) Therefore, I feel the quantum stuff should definitely be deferred to a separate "quantum mechanical explanation of intermolecular interactions" page. OneAhead (talk) 14:31, 22 August 2009 (UTC)

Well I bit the bullet (being bold!) and removed all the QM and put it on a new page. The new page is Quantum mechanical explanation of intermolecular interactions as suggested. Unfortunately, I'm not yet a theoretical chemist (I suppose I'm not technically a chemist at all yet...) so I will have to leave the clean up of the new article to someone who knows it. The stuff I commented last time in this article is still there on the new one. Hopefully someone would maybe be able to do a short intro on this page for the QM theories? I'll have to leave you lot to that! Thecurran91 (talk) 19:23, 30 August 2009 (UTC)

more comments[edit]

The article still needs work - yes, I know, I should get to it! And when I have time I will! There's some discussion of the relative strengths of the different types on the London dispersion forces discussion page:

Peter Atkins and Julio de Paula, pg 703, Atkins' Physical Chemistry, Seventh Edition, 2002, Oxford University Press, "The dispersion interaction generally dominates all the interactions between molecules other than hydrogen bonds."

Ray Eston Smith Jr (talk) 18:50, 15 September 2009 (UTC)

Yep - your last quote (Atkins & de Paula) puts it correctly - unfortunately there is a lot of misapprehension about dispersion forces because the idea that they are weak settled into textbook and teaching mainstream long ago. So editing is needed.

Ian (talk) 11:33, 17 September 2009 (UTC)

As regards what should be in the article - yes it makes sense to have simple descriptions in this article and elsewhere (or later) for more detailed. However, the math description of the perturbation theory derivation is not an explanation, only a math derivation of a result! and the classical/conventional description, as Richard Feynman pointed out (Phys. Rev. 1939, 56, 340), is merely an attempt to interpret the perturbation theory derivation into real terms. What is happening is much simpler - that the electron density of neighboring molecules are being polarized towards each other - this is well established from molecular calculations. This is potentially a static effect (in solids) but will obviously vary in gases or liquids. see R.F.W Bader J. Phys. Chem. A 1998, 102, 7314-7323. But what of this to include as it isn't what is usually found in textbooks?

Ian (talk) 11:33, 17 September 2009 (UTC)

Table of energies way off - units error?[edit]

The table previously described the typical covalent bond energy as 400 kcal/mol, which appears to be off by a factor of 4. Similar factor error observed for the strength of the hydrogen bond, so I changed the units to kJ/mol, which would be roughly correct (and uses SI units) Dllahr —Preceding undated comment added 16:32, 11 November 2009 (UTC).

Someone appears to have changed the units again. Now it is in kJ. I think this is because readers and at least one editor expect to find dissociation energies per bond in the table. Maybe dissociation energies should be listed per bond in eV? At the very least the unit should be changed back to kJ/mol. Maybe it's not good to have such a table comparing intermolecular forces to chemical bonds at all. They are only comparable for solids not liquids and gases where intermolecular bonding changes all the time. (talk) 22:07, 10 December 2009 (UTC)

Looked up the reference Organic Chemistry by Ege (second edition) and found 30-100 kcal/mol for covalent, 3-10 kcal/mol for hydrogen bonds, 1-3 kcal/mol for dipole-dipole interactions, there is no number in there for van der Waals forces. (talk) 10:40, 2 January 2010 (UTC)

This is all wrong. Energy of an H-bond can be as small as zero (obviously), and dipole-dipole interaction (fixed dipoles!) can be repulsive, that is with positive energy.Biophys (talk) 20:28, 2 October 2010 (UTC)

There is a french translation for this page (I don't know how to add it in, will someone take care of this?)

As of 2/12/2010 the units have reverted to kcal. This is surely wrong? Would have thought that kJ/mol would be sensible but I'm not a physical chemist. Somebody competent please. (talk) 13:47, 2 December 2010 (UTC) (talk) 21:36, 9 October 2010 (UTC)

HCl bond[edit]

The bond in HCl is ionic, unequal sharing of electrons, not covalent. — Preceding unsigned comment added by (talk) 14:58, 9 September 2012 (UTC)

not easy to read and memorize[edit]

We need a table to summarize all the intermolecular interactions. You start by mentioning four types, but later in the article you mention hydrogen bond, which you haven't included in the first four types. I understand that they are a subcategory and it may be logical that hydrogen bond wasnt mentioned at the beginning, but this is what I am saying: we need a complete gathered reference of all intermolecular interactions and their characteristics in an one single table. Thanks. — Preceding unsigned comment added by (talk) 11:01, 1 November 2012 (UTC)

London Dispersion[edit]

This second paragraph in the London Dispersion section seems nonsensical and unrelated to dispersion. I will remove it unless anyone knows what this paragraph is meant to say about dispersion (or intramolecular forces in general).

Intramolecular multiple force theory (IMMFT) is a concept which deals with the forces which are developed in different domains of similar supramolecules. The supramolecules such as Dendrimer and others have different discrete zones with their own environment and center of forces and these forces when coordinated to each other in similar molecules, unique molecular mechanics and dynamics are developed known as the intramolecular multiple force theory.[citation needed]

--BBUCommander (talk) 15:16, 11 June 2013 (UTC)

I agree. Do away with it, it makes not sense. kml (talk) 13:26, 18 June 2013 (UTC)

Chronology of inquiry[edit]

Some data concerning the timeline of inquiry of intermolecular forces lack from the article.-- (talk) 15:10, 30 August 2013 (UTC)

Identification of intermolecular force[edit]

The article does not specify how the info on intermolecular force is obtained, through what means.-- (talk) 14:32, 10 December 2013 (UTC)

Effect on the behavior of gases[edit]

I have inserted a new section discussing the effect of intermolecular forces on gases. The article said nothing about why intermolecular forces matter, other than an occasional passing reference to condensed phases and adsorption. I was led to this by making some revisions to the articles on the Joule-Thomson effect and compressibility factor, the latter will be linked to this section. In general, I think that much more could be said about the effects of intermolecular forces. Retired Pchem Prof (talk) 20:39, 16 January 2016 (UTC)

Assessment comment[edit]

The comment(s) below were originally left at Talk:Intermolecular force/Comments, and are posted here for posterity. Following several discussions in past years, these subpages are now deprecated. The comments may be irrelevant or outdated; if so, please feel free to remove this section.

Last edited at 18:43, 4 June 2009 (UTC). Substituted at 19:00, 29 April 2016 (UTC)