Talk:Hydrogen bond

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Contents 1 Hydrogen Bond strength 2 SI-units 3 Lone pairs and HBr 4 Proposal for Clarification 5 Add relative hydrogen bond strengths 6 Diagram is wrong 7 Consistent Dimensions 8 I think there is an error on this page! 9 References 10 Liquid water? 11 Quick Edit 12 I don't understand, is it right? 13 Water Beading 14 -CCl3

[edit] Hydrogen Bond strength

• ) What is the source for the values of the hydrogen bond stregnths (enthalpies)? The accepted value for the H2O...H2O in liquid water is ~ 11 kJ/mol, as was measured by Monosmith & Walrafen (see J. Chem. Phys. 81, 669 (1984)) using Raman spectroscopy.
  

You'll be surprised, but despite its involvment in many water based reactions and biological reactions there is no data regarding the hydrogen bond strengths around the proton! I have added data from a recent article by Markovitch & Agmon that has directed this issue for the first time in liquid water. omermar 23/03/07

[edit] SI-units

Plz, use SI-units! Convert the kcal values kJ.

done, using 1kcal = 4.186kJ then rounding to nearest whole number. Hopefully this is appropriate; frankly, I think those values ought to be sourced anyhow, and will generally vary quite a lot (according to the ionic strength of the medium, and so on) Philbradley 00:58, 12 November 2006 (UTC)

[edit] Lone pairs and HBr

I think most people tend to think of charge as not able to be less than the fundamental electronic charge, but as often much more (as in "Do not enter! High Voltage!"). So I think it's misleading and/or unhelpful to call partial charges "strong" without mentioning that they are indeed partial or providing other guidance as to how strong "strong" is.

Doesn't HBr have three lone pairs of electrons, not two? I don't find the article's description of why HBr has weaker hydrogen bonding that compelling; theoretically an individual HBr molecule could form four hydrogen bonds, right? The catch is that in a pure HBr solution, there are only as many H molecules as HBr molecules so the total number of bonds to be formed is limited by the H molecules, leaving each molecule with a total of two...

A hydrogen bond only forms when the H atom is attached to either an F,O, or N atom, and is "sufficiently close" to another H atom also attached to an F,O or N atom. The reason is because the F, O and N atoms are very electronegative and their nuclei are very small and (and so the charge density is relatively high). The high electronegativity atom draws the electron away from the hydrogen atom, leaving the proton relatively exposed. The proton (slightly positive in charge) can now approach an electronegative atom, and form the "hydrogen bond". It has to approach an F,O or N atom, because (you can think of it this way) their nuclei are about the same size as the exposed proton, which means the hydrogen bond formed will be a good fit. This is sort of a hand-wavy explanation, but the essentials are all here. Yes, the hydrogen atom does participate in other intermolecular forces, but they aren't anywhere as strong as what we typically think as "hydrogen bonding" - as such, they aren't as important. HappyCamper 05:17, 28 Mar 2005 (UTC)

• Also, the Lewis dot diagram for HBr does indeed indicate that Br has 3 lone pairs. However, it is important to recognize that these "lone pairs" are only used as a heuristic to understand the organizational structure that is at the heart of chemistry. In fact, quantum mechanical calculations have shown (for example, the water molecule), there are actually no "bunny ears" sticking out from the O atom which are often used to represent the two lone pairs there. The chemistry of water just behaves as if it did, and so using lone pairs to designate this should be understood as strictly a tool. HappyCamper 05:17, 28 Mar 2005 (UTC)
• The "catch" that you mentioned has doesn't have much to do with the lack of hydrogen bonding for HBr. It has to do with the fact that the Br atom is very big. Even if the proton from H can hydrogen bond to it, the charge would be spread out over such a large area that the resulting bond would be very very weak. Keep in mind, however that the hydrogen bonds will form and break very frequently if the temperature is high enough. You might be interested to know, for example, that HF forms hydrogen bonds, and in fact, it is possible for HF to form rings of 5 molecules, all bonded together with hydrogen bonds! Granted, the bonds will break and spontaneously form other structures. HappyCamper 05:17, 28 Mar 2005 (UTC)
• And yes, I agree with your first paragraph, but I think in this context it isn't necessary to introduce the complication behind how electronegativities are derived. HappyCamper 05:17, 28 Mar 2005 (UTC)

[edit] Proposal for Clarification

I think we should mention on the page somewhere that the hydrogen bond is not necessarily intermolecular. It can be intramolecular as well. Consider the compound H2NCH2CH2CHO for example (1-aminopropanal). The H atom attached to the N atom can hydrogen bond to the aldehyde end! HappyCamper 05:17, 28 Mar 2005 (UTC)

True - proteins are good example as well (perhaps the reader may connect more redily with this example and there is plenty on the web about these H-bonds).

[edit] Add relative hydrogen bond strengths

Can someone look up in a table the range relative strengths of hydrogen bonds in these configurations? HappyCamper 05:17, 28 Mar 2005 (UTC)

  F-H ..... H-F
  F-H ..... H-O-R
  F-H ..... H-N-R
R-O-H ..... H-O-R
R-O-H ..... H-N-R
R-N-H ..... H-N-R

[edit] Diagram is wrong

The picture for this article shows water molecules all in a jumble. Hydrogen bonding does not allow the water to bunch up in the patterns in the picture - it prefers that the hydrogen lie on the straight line drawn between the two heteroatoms. Basically, the H-bond angles (the angle from O to H to O) should be close to 180˚ as possible. If a Hydrogen bond angle deviates from 180˚ by more than 30˚, the strength of the bond goes to zero and the hydrogen bond disintegrates. (this is how water evaporates)

What do you think of this diagram? --JWSchmidt 00:18, 6 October 2005 (UTC)

I think the diagram in this aricle is fine because it shows dynamic nature of "transient" H-bonds in liquids. The H-bonds are more perfect in solids, but the actual distributions of such angles in molecular crystals are rather broad, and their maxima are sometimes not 180 degrees. Biophys 06:00, 6 November 2006 (UTC)

The molecules are in a jumble because this is a picture of water molecules in the liquid phase, which is naturally "jumbled". Hydrogen bonds are much less "rigid" than people think. Even if the minimum energy configuration has an angle of 180 deg, the molecules are constantly rotating and translating (especially in the liquid phase), which results in a broad distribution of angles. The higher the temperature, the less "perfect" the bonding pattern, due to the increasing importance of entropy. -- Itub 14:41, 7 November 2006 (UTC)

[edit] Consistent Dimensions

In the introduction, in different contexts, two different units are given for H-bond energies, kcal/mol and kJ/mol. Obviously it's better if only one unit is adopted, or values are given in both units. 99of9 23:52, 7 November 2005 (UTC)

[edit] I think there is an error on this page!

I could be very much mistaken but I believe the following data in the main article is wrong:

O—H...:N (7 kcal/mol) O—H...:O (5 kcal/mol) N—H...:N (3 kcal/mol) N—H...:O (2 kcal/mol)

Should it not read as follows?

O—H...:N (7 kJ/mol) O—H...:O (5 kJ/mol) N—H...:N (3 kJ/mol) N—H...:O (2 kJ/mol)

---

Please provide a source. --JWSchmidt 13:55, 6 December 2005 (UTC)

I can't find the original source where I got the initial 4 values, so I rechecked two of them. For NH₃ (N-H...:N), the value is about 3.3kcal/mol (Solomons, T.W. Graham (1988). Organic Chemistry, 4th Ed. John Wiley & Sons, p88). But for H₂O (O-H...:O), the value seems to vary considerably. Some published values are 8.7 kcal/mol (Solomons, 1988) and 5.58 kcal/mol (Suresh, S.J., Naik V.M. "Hydrogen bond thermodynamic properties of water from dielectric constant data", J. of Chemical Physics, 1 Dec 2000, 113, 21). Some other H-OH...OH₂ energy values from the internet are: 4.7-5 kcal/mol [2], 6.6 kcal/mol [3]), and 6 kcal/mol [4]. According to the paper by Suresh a wide range of values from 3-8 kcal/mol have been reported, and different techniques are used (e.g. IR absorption, NMR shift, X-ray, Neutron diffraction). Hence the value of 5 kcal/mol stated in the original reference is probably just an approximate value. (Note: 1 kcal/mol = 4.1868 kJ/mol) Nathaniel 07:45, 7 December 2005 (UTC)

Thanks for checking into this. I'm not a chemist, but it makes sense to me that it would be hard to measure the energy and that it would be dependent on conditions during the experiment. --JWSchmidt 17:48, 7 December 2005 (UTC)

But they are still much weaker than covalent bonds. Biophys 06:02, 6 November 2006 (UTC)

The estimates of H-bond energy should be described more carefully. First, is it free energy or enthalpy? Second, energies of H-bonds are different in vacuum and in different media. Third, energies of H-bonds could correspond either to enthalpy of sublimation or enthalpy of fusion. Biophys 06:36, 6 November 2006 (UTC)

[edit] References

I suggest adding some external links from this page. Semi Psi 01:15, 3 February 2006 (UTC)

[edit] Liquid water?

In the caption to the picture, how about "liquid H2O" or just "water"?

01:47, 9 May 2006 (UTC)

[edit] Quick Edit

I have removed some 11-year-old kid rants ("lalalala" or something) from the beginning of this article. This is the first time I have to edit a page of Wikipedia... And I liked it!. Maybe I´ll give a hand, as a registered user :D

[edit] I don't understand, is it right?

I am citing from the text: --oxygen, nitrogen or fluorine, are the doners!!!!!???? Or the receivers of the electron!! --This electronegative element attracts the electron cloud from around the hydrogen nucleus and, by decentralizing the cloud, leaves the atom with a !!!positive!!!???? partial charge. --hydrogen bond results when this strong ?positive? charge density attracts a lone pair of electrons on another heteroatom, which becomes the hydrogen-bond acceptor. !?If it attracts should it be the acceptor?! Now it say oxygen, nitrogen or fluorine are the aceptors?! (i think this is rigth) It seams to me that since the electronegative of hidrogen is samaller then the electronegative of oxygen, nitrogen or fluorine then hidrogen is the doner...--Paclopes 22:19, 24 July 2006 (UTC)

The donor is the atom that "donates" the hydrogen atom. For example, in ROH ... NR3, the alcohol on the left is donating a hydrogen bond to the amine on the right. Itub 12:36, 25 July 2006 (UTC)

Yes. I'd like to extend this answer and say that in H2O, the oxygen can either act as a DONOR by donating a hydrogen to form a hydrogen bond with another molecule, but it could very well act as an ACCEPTOR when one of its lone pairs (pairs of electron not participating in a covalent bond) accepts a hydrogen from another molecule. omermar 24/03/07

[edit] Water Beading

I think in the water section someone should add that hydrogen bonding is the reason that water beads and does not just stay flat when spilled or put on a flat surface. All the water molecules are weakly bonded to eachother and can easily be pushed apart if you push the bead flat.Kniesten 18:07, 6 September 2006 (UTC)

[edit] -CCl3

Not only just N O and F can form hydrogen bond but also -CCl3

For example, chloroform has hydrogen bond.

I've added this information to the article, but remember that you can edit it yourself. This is a wiki, after all! --Itub 12:53, 20 October 2006 (UTC)