Talk:Intermolecular force

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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.

Contents

[edit] Van der Wall's information incorrect?

  • 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)

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

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)

[edit] Van der Waal's Volume?

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.

83.103.67.194 13:26, 28 March 2007 (UTC)roberto90967

[edit] Diagrams

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

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

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 68.63.61.216 (talk • contribs) .

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

[edit] Polar molecules not in Symmetrical molecules

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 82.22.102.86 (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)

[edit] Intro

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)

[edit] Merging Keesom force to Intermolecular Force

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

[edit] clean up

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.

[edit] Possible mistake

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 203.87.191.146 (talk) 10:52, 3 September 2007 (UTC)

[edit] This article is terribly useless.

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.--74.61.4.8 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)

[edit] Overall critic

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 (talk • contribs) 21:09, 25 May 2008 (UTC)