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I'm not satisfied with the rationale for the intermolecular attraction known as hydrogen bonding. In my book, it states that

Hydrogen bonding is a special type of intermolecular attraction between the hydrogen atom in a polar bond (particularly H ¬ F, H ¬ O, and H ¬ N) and non-bonding electron pair on a nearby small electronegative ion or atom usually F, O, or N (in another molecule).

It seems that chemists looked at the data and found they needed a 'fudge factor' to fit the higher boiling points of ammonia, water, and hydrogen fluoride.

Why doesn't hydrogen bonding apply to other atoms like Sulfur or Chlorine? They seem to be electronegative enough IMO. Is there an explanation that explains why this is, rather than saying that it's there? What makes Hydrogen so special?

Note: Chlorine is more electronegative than Nitrogen.

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    $\begingroup$ Hydrogen bonds are observed also for heavier atoms, but are generally weaker. $\endgroup$ – Mithoron Mar 5 '15 at 21:54
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    $\begingroup$ One thing to note: Chlorine is less electronegative than nitrogen. The $\ce{NCl3}$ molecule, when reacting with a chloride source releases $\ce{Cl2}$ showing that the chlorines in $\ce{NCl3}$ have a positive partial charge. $\endgroup$ – Jan May 17 '15 at 20:07
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    $\begingroup$ @Jan Wikipedia says that the Pauling EN of Cl is more than N. $\endgroup$ – user223679 May 18 '15 at 4:16
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    $\begingroup$ @user223679 And right underneath it says that Cl is less electronegative than N in the Allen scale. There are many electronegativity scales out there. I choose to believe the experiment, if it contradicts the theory. $\endgroup$ – Jan May 19 '15 at 21:48
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There are two main ways to look at hydrogen bonding. The first is electrostatic, where the electronegativity of the atoms is used to describe the interaction. Your argument about chlorine being more electronegative than nitrogen is a good one suggesting that the electrostatic argument is only part of the story, and there is at least one study that suggests that hydrogen bonding does occur with chlorine in polar molecules such as chloroform.

We can increase dramatically the number of atoms considered to undergo hydrogen bonding if we take a molecular orbital approach. My reasoning here is a summary of what can be found in Inorganic Chemistry by Miessler, Fischer and Tarr. They, in turn, rely heavily on a new definition of hydrogen bonding recommended by the IUPAC Physical and Biophysical Chemistry Division.

A hydrogen bond is formed when an $\ce{X-H}$ (where X is more electronegative than H) interacts with a donor atom, $\ce{B}$. The attraction $\ce{X-H...B}$ can be described as consisting of the following components:

  • An electrostatic contribution based on the polarity of $\ce{X-H}$
  • A partial covalent character which arises from the donor-acceptor nature of the interaction
  • Dispersion forces

The first bullet is typically the only phenomenon discussed in General Chemistry classes, and this is not unreasonable since the second bullet requires the introduction of Lewis Acid/Base concepts which may not have been covered.

It is important to note that the new definition includes what I like to call "the proof is in the pudding" where the existence of hydrogen bonding requires experimental evidence, which can be found using a number of methods:

  • The $\ce{X-H...B}$ bond angle: a bond angle of 180 degrees indicates strong hydrogen bonding and would be accompanied by shord $\ce{H...B}$ bond distances.
  • A red shift in the IR frequency upon formation of $\ce{X-H...B}$
  • Hydrogen bonding results in deshielding of the H atom, which can be observed in NMR spectroscopy.
  • Thermodynamic studies should indicate that $\Delta G$ for the formation of $\ce{X-H...B}$ should be larger than the thermal energy of the system.

While electrostatics is the dominant contributor to hydrogen bonding, an analysis of frontier orbitals can also be insightful. A thorough answer requires a fair amount of MO theory background, but for those interested, looking at the MO diagram of $\ce{FHF-}$ will be helpful. In brief, the base ($\ce{F-}$ in this case) has a filled $p_z$ orbital that gains access to a delocalized orbital with lower energy, thus validating the formation of a strong hydrogen bond.

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    $\begingroup$ Is there a reason as to why hydrogen bonds are stronger than other dipole-dipole interactions? $\endgroup$ – 1110101001 Mar 7 '14 at 4:47
  • $\begingroup$ This website (and also the definition the OP gave) mention "electron lone pairs" as an important criterion for the Hbond to be formed. Please share your comment on that. Thank you! $\endgroup$ – Gaurang Tandon Jan 3 '18 at 16:13
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N, O, F have atomic numbers $Z=7, 8, 9$, respectively. Chlorine has $Z = 17$ which is much larger. Consequently, the atom is larger as well and it is more diffuse. Equivalently, the lone pairs in chlorine are at the 3-level which is too high. Because the hydrogen bond isn't a real bond but a dipole-dipole attraction and because the force between two dipoles scales like $1/r^4$, a larger atom implies a significantly weaker dipole-dipole attraction which is, in the case of chlorine, beaten by other forces. That's why hydrogen bonds are only observed with light atoms bound to hydrogen.

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    $\begingroup$ There could be a molecular orbital aspect to hydrogen bonding, see @bob's answer. But this mode, too, favors more compact atoms like N instead of Cl. $\endgroup$ – Oscar Lanzi Mar 25 '17 at 22:37

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