Talk:Hydrogen bond

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[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 [1], 6.6 kcal/mol [2]), and 6 kcal/mol [3]. 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)

This is a strange way of saying things. Yes the ultimate cause is H-bonding but the usual explanation is surface tension(, due to the cohesion of water, due to H-bonding). -User: Nightvid

[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)


[edit] Theory, covalent nature, references

About this : "The initial theory of hydrogen bonding proposed by Linus Pauling suggested that the hydrogen bonds had a partial covalent nature. This remained a controversial conclusion until the late 1990's when NMR techniques were employed by F. Cordier et al. to transfer information between hydrogen-bonded nuclei, a feat that would only be possible if the hydrogen bond contained some covalent character." and the reference given (Cordier et al., J. Magn Res. 1999, 140, 510-512): I have just looked at this article and it is clearly written in it that other examples of J coupling through hydrogen bonds were known at least one year before them. The most ancient article about this phenomenon seems to be : A. J. Dingley and S. Grzesiek, J. Am. Chem. Soc. 120, 8293–8297 (1998). I hesitated to modify the article myself as I am not sure to understand: is the observation of such a coupling an evidence of the covalent nature of the hydrogen bond ? Or is there something new and decisive on this question in the article of Cordier et al. ? I wonder because they don't claim it at all in their article...

Other articles that might be relevant (I haven't read any of them, but they were cited in Weinhold and Landis's Valency and Bonding, Cambridge University Press, 2005):
  • Summers, MF. J Am Chem Soc 114, 4391 (1992)
  • Wütrich K. Proc Natl Acad Sci USA 95, 14147 (1998)
  • Shenderovich, SN et al. Ber Bunsenges Phys Chem 102, 422 (1998)
  • Cornilescu JS. J Am Chem Soc 121, 2949 (1999)
  • Wang YX. J Biomol NMR 14, 181 (1999)
There have also been discoveries of long-range quantum-mechanical phase coherence that have been used as evidence of the covalent character of the H-bond:
  • Isaacs ED et al. Phys Rev Lett 82, 600 (1999); Science 283 (1999)
I think it would be best not to try to attribute the discovery of evidence of covalent character in hydrogen bonds to any specific author in the article text, because it can be controversial. We can cite some of these references, of course, or a more indirect source such as Valency and Bonding (p. 583). --Itub 09:54, 3 April 2007 (UTC)

[edit] Hydrogen Bond Angles

Would it be feasible to add some more discussion of typical donor-hydrogen-acceptor angles? I know that there is often significant variation, however many geometric analysis software packages such as VMD require a user to input distance and angle values to determine the existence of hydrogen bonds given a set of molecular positions such as a box of water or a protein structure. A few brief comments in this regard would be very helpful. Example: typical donor-acceptor distances are near 3+/-0.2 A with donor-hydrogen-acceptor angles near 180+/-30 degrees.
Paul.raymond.brenner 16:30, 25 April 2007 (UTC)

You are welcome to add them, just try to add a reference for the values if possible. --Itub 08:21, 26 April 2007 (UTC)
You may recieve some analysis I did on this point. Visit my page (www.fh.huji.ac.il/~omerm) to get my email. Omermar @ 28/2/2008 —Preceding unsigned comment added by Omermar (talk • contribs) 17:39, 28 February 2008 (UTC)

[edit] "Carbon can form hydrogen bonds" <- disputed

"Chemistry", Gregory M. Williams, John A. Olmsted.

"Carbon never forms hydrogen bonds". Contributer314 12:00, 27 May 2007 (UTC)

This is a generalization on the part of the book. Carbon usually doesn't participate in H-bond formation. However, the dissociation of the acid HCN in water shows that it must be happening. -User: Nightvid

[edit] Why not chlorine?

My table shows nitrogen and chlorine both having electronegativities of 3.0 - so why can't a H-Cl bond produce hydrogen bonding? Something to do with a larger atomic radius from being in the higher period? Jasonfahy 19:06, 29 May 2007 (UTC)

The lone pairs of electrons on the Cl are in the third shell, unlike N, O, and F. They are thus spread out more and don't have a concentrated negative charge density, which is necessary for H-bond formation. -User: Nightvid

[edit] Hidden content

There's some "hidden" content in the Advanced theory of the hydrogen bond section, enclosed in comment <!-- --> tags. Why were these texts removed? --Freiddie 19:59, 5 August 2007 (UTC)

  • I commented that text out, because I don't understand any of it (and I have close to 40 years of experience in research in intermolecular forces). I did not remove the text altogether, because of the chance that the original author can explain what it means. In that case we can reinsert some of this text after some discussion and clarification. However, you are the first one to react. I assume you don't know either what a metric dependent electrostatic scalar field between two or more intermolecular bonds is and what its relation is to hydrogen bonding? --P.wormer 12:32, 7 August 2007 (UTC)
I hardly know what that statement means when all those "slightly understood" words are put together in one phrase. I'm merely puzzled by the extra spaces in the text at the end of the section (I thought it was just some extra blank space until I edited it to remove it, when I realized that it was not a blank space after all). --Freiddie 15:52, 8 August 2007 (UTC)

[edit] Sniffle wiffle?

It says that hydrogen bonding is called "sniffle-wiffle" (do Ctrl-F to find it). What does that mean? Does anyone have a source? Shalom (Hello • Peace) 00:45, 12 October 2007 (UTC)

That was vandalism, from here. But this makes me very wary about how much hidden vandalism there is in many of hour pages. --Rifleman 82 04:38, 12 October 2007 (UTC)

[edit] Another ?

It says in the article:

  • HO—H...:OH3+ (18 kJ/mol[1] or 4.3 kcal/mol) {Data obtained using molecular dynamics as detailed in the reference and should be compared to 7.9 kJ/mol, obtained using the same molecular dynamics.}

Could someone please explain what this means? Shalom (Hello • Peace) 00:51, 12 October 2007 (UTC)

  • Yes. "Bulk" water hbond (that is - hydrogen bond between two regular water molecules in pure water) enthalpy is about 11 kJ/mol, that is - you need to apply 11 kJ of energy to break 1 mol of hbonds. Zero point energy accounts for another 11 kJ/mol. This might cause some confusion - some people prefer to report the 11 kJ/mol value while other prefer the 11*2=22 kJ/mol value.

In simulations, one often makes some assumptions to simplify the problem and make it possible to model it. For water, many simulations describe a water molecule as being made out of 3 point charges (1 to represent the Oxygen and 2 more to represent the hydrogens). This is, ofcourse, a simplified picture becuase water molecule is actualy a continum of charge density. When you make such assumptions some of the values you calculate are usualy a bit off compared to the values measured in experiments. This is the reason why the 7.9 kJ/mol is the value of bulk hbond, and not 11 kJ/mol. This point needed to be stressed out because the 18 kJ/mol is about 60% higher then 11 kJ/mol but 225% higher then 7.9 kJ/mol. I hope this wasn't too long of an answer, omer. 30/01/2008 omermar —Preceding comment was added at 13:25, 30 January 2008 (UTC)