Normally, chlorine doesn't form hydrogen bonds because despite its electronegativity, the size of the atom is such that its electron density is too low to form hydrogen bonds.

However, chlorine forms hydrogen bonds in chloral hydrate (2,2,2-trichloroethane-1,1-diol). What could be the possible reason for this?

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  • $\begingroup$ I am far from to be a specialist of organo-chemistry. However I remember learning that geminal diols are particularly unstable. So I would say maybe in this case the hydrogen atoms exhibit "a smaller electronic density" (the electrons are diffused in the whole molecular orbitals) then even the small electronic density of the chlorine atoms would be enough to get a very weak hydrogen bond. $\endgroup$
    – ParaH2
    Commented Jun 19, 2018 at 18:59
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    $\begingroup$ Possible duplicate of Is hydrogen bonding generally defined to include only three period two elements? $\endgroup$
    – Mithoron
    Commented Jun 19, 2018 at 19:38
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    $\begingroup$ In other words, your starting premise is false - Cl hasn't got much problem with hydrogen bonding. It's just weak. For some reason weak hydrogen bonds often aren't considered at all. $\endgroup$
    – Mithoron
    Commented Jun 19, 2018 at 19:41
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    $\begingroup$ How is that a duplicate if chloral hydrate includes a $\ce{Cl\bond{...}H}$ hydrogen bond? Or is this a misconception? $\endgroup$ Commented Jun 19, 2018 at 20:36
  • $\begingroup$ @pentavalentcarbon it is misconception to be debunked. I couldn't find straight up dupe, but there as many posts about it. $\endgroup$
    – Mithoron
    Commented Jun 20, 2018 at 16:53

2 Answers 2


As noted in the comments, the fallacy is the notion that only a select few atoms can form hydrogen bonds with protic hydrogen. In reality, the strongest hydrogen bonds involve nitrogen, oxygen, or fluorine, but most nonmetals can form such bonds if they have an electron pair to donate to the bond. Hydrogen bonding with chlorine is seen in the $\ce{HCl_2^-}$[1], which is analogous to the more familiar bifluoride ion, as well as chlorine-bearing organic compounds[2].


1. Harry F. Herbrandson, Richard T. Dickerson Jr., and Julius Weinstein, " Cite this: J. AThe Bichloride Ion", J. Am. Chem. Soc. 1954, 76, 15, 4046. https://doi.org/10.1021/ja01644a066

2. R. Banerjee, G. R. Desiraju, R. Mondal, J. A. Howard. "Organic chlorine as a hydrogen-bridge acceptor: evidence for the existence of intramolecular O--H...Cl--C interactions in some gem-alkynols." Chemistry 2004 Jul 19; 10(14):3373-83. https://doi.org/10.1002/chem.200400003. PMID: 15252783.


In chemistry, generally most of the notable exceptions are associated with the thermodynamics or the kinetics of the reaction.

In the particular example of chloral hydrate, chlorine tends to form hydrogen bonds with hydrogen (even though usually their hydrogen bonds don't exist) majorly due to the thermodynamic changes that occur in the process. By forming such "pseudo-rings" (because H-Bonds aren't completely a "bond", rather only very strong attraction) the structure of chloral hydrate gains excessive stability as 5 & 6 membered rings are very stable. This pseudo-ring formation is called "Chelation Effect".

So, basically that extra stability is the driving force for chlorine to form H-Bonds and the process, as a whole, becomes thermodynamically feasible (spontaneous).


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