I understand that dipole dipole forces is due to the attraction of the different partials charges of atoms in different molecules due to their different electro-negativities.

For hydrogen bonding, my understanding of it is that they are just a special case of dipole dipole forces. In certain molecules in where the difference in electro-negativities between the different atoms is high enough, then even stronger dipole dipole forces will be form which are known as hydrogen bonding. For example, in $\ce{H_2O}$ where the difference in electro-negativity of H and O is 1.24, hydrogen bonding occur. This is why hydrogen bonding can only occur between hydrogen and nitrogen, fluorine or oxygen.

However, my school teacher says that hydrogen bonding in fact is not a type of dipole dipole force as she says that hydrogen bonding occurs differently to dipole dipole forces. She said that hydrogen bonding occurs to the attraction of the lone pairs in oxygen, fluorine and nitrogen atoms to atoms with a positive partial charge. Hence this is why hydrogen bonding is much stronger than conventional dipole dipole forces.

But this doesn't make sense because if hydrogen bonding is due to the attraction of H to the lone pair of electrons, shouldn't hydrogen bonding also form between H and Cl or I or any other atom that has a lone pair and is more electronegative than H?

Edit 1

From the comments below, it appears that the actual answer is more complicated than the simple explanation of hydrogen bonding given above. After doing some research, I have became aware of the Advanced Theory of the Hydrogen Bond which states that the hydrogen bond actually contains 10% covalent nature. However none actually explain how they are formed. Could someone please explain how hydrogen bonding actually in reality form.

Edit 2

So in response to the excellent answer by jheindel, is the reason why hydrogen bonding is considered partly covalent in nature is because the let say in water, oxygen's MO partly overlaps with the MO of the hydrogen atom in the other molecule. Also this is why they are much stronger than conventional dipole dipole forces?

  • $\begingroup$ @Martin -If you don't have time to answer, could you please tell me what area of chemistry the answer to my question involves. Also if there are any, is there any links you could give me that explain my question. Sorry if I sound really annoying and desperate, but I really want to know and I have been scanning the internet for days for an answer and I couldn't find anything that was remotely close to answering my question. $\endgroup$
    – Nanoputian
    Commented Aug 28, 2015 at 12:32
  • $\begingroup$ related chemistry.stackexchange.com/questions/24157/… $\endgroup$
    – Mithoron
    Commented Aug 29, 2015 at 13:23
  • $\begingroup$ I hope this will help you. adichemistry.com/general/chemicalbond/hbond/… $\endgroup$ Commented Jun 2, 2016 at 7:45
  • $\begingroup$ Here is some experimental and computational data on very strong hydrogen bonds with covalent character: science.sciencemag.org/content/371/6525/160 $\endgroup$
    – Karsten
    Commented Jan 13, 2021 at 16:41

3 Answers 3


Well, it turns out that this is a very active area of research. I will only summarize what I understand to be true about the covalent nature of the hydrogen bond, so I'm sure the explanation could be more detailed and potentially more accurate in some places (I hope someone gives a more detailed answer), but here's what I've got.

As you said, it has been found that the hydrogen bond (specifically in water) is about 90% electrostatic and 10% covalent. See this link for a sort-of-review/overview of hydrogen bonding and some relevant research.

Upon following the paper trail, I found the following research article: E.D. Isaacs, A. Shukla, P.M. Platzman, D.R. Hamann, B. Barbiellini, C.A. Tulk, J. Phys. Chem. Solids, 2000, 61, 403-406. (Mirror)

Essentially, they use some method which I don't really understand and find quite conclusive evidence of the (word of the day) anisotropy of hydrogen bonds. That means that the bond is indeed directionally dependent. Bond direction is one defining characteristic of covalent bonds.

That raises the question of what is going on with the orbitals that allow for hydrogen to participate in more than one covalent bond at a time.

A simple explanation of the orbital interactions for a water dimer is to view the oxygen 2p,2s hybridized orbital as having superposition with the hydrogen 1s orbital and then interacting covalently with that same 2p,2s hybridized orbital on the other oxygen in the dimer. This superposition which includes the hydrogen then explains the way that hydrogen is participating in one covalent bond and another partially covalent interaction.

Furthermore, this explanation of what the orbitals are doing fits well with what we would expect from a hydrogen participating in two different covalent interactions, which is that the energy curve has two minima where one is much deeper than the other.

Further evidence for the covalent nature of the bond is the fact that the O-H bond elongation and O---H hydrogen bond length always change together. That is, as the O-H bond lengthens, the O---H interaction length shortens. This implies a shift in electron density towards hydrogen which is most easily explained by a delocalization of electrons participating in a covalent interaction.

So, I hope that's a good start toward understanding the covalent nature of the hydrogen bond--it was for me!

But to answer the simple part of the question which you clearly already know is that yes dipole-dipole is a significant part of hydrogen bonding, and it certainly should be, but as is always true of chemistry, there's a lot more going on than what one might initially think.

EDIT (06/01/2016)

Since I posted this answer, I think I have developed a better understanding of the situation, and feel obliged to update my answer. First, let me split hydrogen-bonding into two obvious parts, electrostatic and covalent characteristics. This is an obvious split, but it is not completely clear yet what we mean by covalent, so let us try to define that by looking at some examples. Furthermore, we must be careful when talking in about hydrogen-bonding because in every hydrogen bond there is a donor and acceptor, and these two roles are often played best by different chemical species.

Hydrogen-Bonding in thiols

Here we go breaking yet another rule. They told us that sulfur couldn't hydrogen bond. Well, that's not exactly true. Okay, but that hydrogen-bonding has got to be basically negligible right? Not exactly.

The first thing we need to do here is get some information. First, sulfur has an electronegativity of 2.58 on the Pauling scale. This is roughly the same as the electronegativity of carbon which we do not consider to hydrogen-bond well. So, if there are hydrogen-bonds involving sulfur, we should expect that they will be primarily covalent as any dipole interactions will be pretty small.

As it turns out, sulfur is a much better hydrogen-bond acceptor than it is donor, and at times is a comparable acceptor to oxygen. Here is a table comparing certain dimers from a great source you can reference at the bottom of the page$^3$.

enter image description here

As you can see, when the sulfur in these dimers (which are all named in the caption of the table) acts a hydrogen bond acceptor, the shift in vibrational frequency of the $\ce{O-H}$ and the binding energy are comparable to when an oxygen is the acceptor atom.

Looking at other research$^4$ involving sulfur hydrogen bonds reveals that approximately 70% of the interaction energy between sulfur and oxygen when sulfur is the acceptor comes from dispersion. That is the key physical insight in all this. Despite the loss of electronegativity, sulfur gains something in its greater dispersion interactions which makes it a good hydrogen bond acceptor. It should be noted, however, that the electronegative component of these hydrogen-bonds maintains their directionality.

Also, to compare, according to that wikipedia article I linked, the interaction energy of a water dimer is only 24% dispersion. In water electrostatics dominate and induction and exchange play meaningful roles.


I have more I could say, but it's perhaps a longer explanation than is worth going into. To conclude, what is meant by covalent interactions in hydrogen-bonding varies based on the species you are considering. In something like water, we can think about this in the usual way by saying that there must be some orbital overlap and hence electronic exchange which will be stabilizing. With other species, we can see that what we mean by covalent interactions might be different. That is, there is still some orbital overlap, but we must also consider non-electrostatic interactions like dispersion.

3 Noncovalent Forces-Chapter 2

4 Wennmohs, F., Staemmler, V., & Schindler, M. (2003). Theoretical investigation of weak hydrogen bonds to sulfur. The Journal of chemical physics, 119(6), 3208-3218.

  • $\begingroup$ I am little confused when you say, ' view the oxygen 2p,2s hybridized orbital as having superposition with the hydrogen 1s orbital and then interacting covalently with that same 2p,2s hybridized orbital on the other oxygen in the dimer'. Do you mean at the end of the sentence it is interacting covalently with the other hydrogen atom or am I missing something here? $\endgroup$
    – Nanoputian
    Commented Oct 8, 2015 at 9:26
  • $\begingroup$ What I mean is that the acceptor oxygen sp3 (if you still want to use that idea) does, to some extent, overlap with the hydrogen 1s of the donor molecule. This is the covalent nature of the bond. Sorry, I definitely worded that poorly. $\endgroup$
    – jheindel
    Commented Oct 8, 2015 at 16:31
  • $\begingroup$ I would also point out that this is actually much more complicated than my answer and I don't feel like I understand all the details well enough to lay them out clearly. Hopefully Martin answers. $\endgroup$
    – jheindel
    Commented Oct 8, 2015 at 16:33

Hydrogen bonding may be understood in different ways. My own preference is it us a form of delocalized covalent bonding. This is discussed in this answer.

This molecular orbital description does not require the bonding to be polar, although the polarity commonly associated with hydrogen bonds (hydrogen primarily attached to a much more electronegative atom in both the primary bond and the hydrogen bond) favors more bond delocalization and thus a stronger hydrogen bond. Answers to this question explore the possibilities, including one where that cites hydrogen bonding to carbon atoms.


Hydrogen bond is defined as the electrostatic force of attraction which exists between the covalently bonded hydrogen atom of one molecule and the electronegative atom of other molecule. These are special type of dipole-dipole interactions. Here the hydrogen atom is linked with one atom through covalent bond and linked with other atom through hydrogen bond.however hydrogen bond is much weaker than the covalent bond. When hydrogen atom bonded to a strongly electronegative atom the bond pair of electrons displaced towards electronegative atom. Thus the only electron of hydrogen lies away from it, its nucleus get exposed and behaves as a bare proton and exert a strong electrostatic force of attraction on the strong electronegative atom of adjacent molecule this interaction is referred as hydrogen bond.

  • $\begingroup$ So what does the lone pair of the oxygen, nitrogen or fluorine atom got to do with hydrogen bonding? $\endgroup$
    – Nanoputian
    Commented Aug 26, 2015 at 9:31
  • $\begingroup$ @Nanoputian - its a sink of negative charge which attracts the somewhat 'bare', positively-charged proton (i.e. the hydrogen involved in H-bonds). $\endgroup$ Commented Aug 26, 2015 at 11:32
  • 4
    $\begingroup$ Sorry, but telling that h-bond is purely electrostatic is a fairytale for little students, told because truth is hard to understand. $\endgroup$
    – Mithoron
    Commented Aug 26, 2015 at 11:46
  • 1
    $\begingroup$ Don't worry about being too complicated for a high school student to understand. I am familiar with 1st and 2nd year university chemistry and should be able understand how h-bonding occurs in reality. $\endgroup$
    – Nanoputian
    Commented Aug 26, 2015 at 12:54
  • 2
    $\begingroup$ This answer is technically wrong, but it's not delete worthy, VLQ, or NAA. Please don't flag it for no reason. $\endgroup$
    – M.A.R.
    Commented Aug 27, 2015 at 13:23

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