Hydrogen bonding and electronegativity

So once again it's #RealTimeChem Week and to 'celebrate' we'll be taking a look at the relationship between hydrogen bond basicity and electronegativity in this blog post. The typical hydrogen bond is an interaction between an electronegative atom and a hydrogen atom that is covalently bonded to an another electronegative atom. We tend to think about hydrogen bonding as electrostatic in nature and we often use electrostatic models to describe the phenomenon. Let's take a look at hydrogen fluoride dimer which is probably the simplest hydrogen bonded system.

Fluorine is more electronegative than hydrogen which means that it tends to draw the electrons it shares with hydrogen towards itself. This gives fluorine a partial negative charge and hydrogen a partial positive charge. This simple electrostatic model suggests that a hydrogen bond will get stronger in response to increases in the electronegativity of either the acceptor atom or the atom to which the donor hydrogen is covalently bonded. 

Hydrogen bond strength can be quantified as the equilibrium constant for the association of a hydrogen bond donor with a hydrogen bond acceptor in a non-polar solvent. For example, pK_BHX can be used as a measure of hydrogen bond basicity where K_BHX is the equilibrium constant for association of the hydrogen bond acceptor compound (e.g. pyridine) with 4-fluorophenol in carbon tetrachloride.  Let's take a look at some pK_BHX values for three structurally prototypical  compounds that present nitrogen, oxygen or fluorine to a hydrogen bond donor. 

The trend is the complete opposite of what you might have expected on the basis of the simple electrostatic model for hydrogen fluoride dimer. However, this is not as weird as you might think because electronegativity tells us about distribution of charge between atoms but at hydrogen bonding distances the donor can 'sense' the distribution of charge within the acceptor atom. Electronegativity quantifies the extent to which an atom can function as an 'electron sink' and this is also related to how effectively the atom can 'hide' the resulting excess charge from the environment around it. Put another way, fluorine will appear to be really weird if you think of it as a large, negative partial atomic charge

This is a good place to wrap up and, if you're interested in this sort of thing, why not take a look at this article on prediction of hydrogen bond basicity from molecular electrostatic potential. My most up to date hydrogen bond basicity data set can be found in the supplemental information (check the zip file) for this article and that's where I got the figures for the table.