For a hydrogen atom in a molecule to participate in a hydrogen bond, several conditions must be met beyond just the presence of hydrogen:
- The hydrogen atom must be covalently bonded to a highly electronegative atom such as nitrogen (N), oxygen (O), or fluorine (F). This bond is highly polar because the electronegative atom pulls electron density away from hydrogen, creating a partial positive charge (δ+) on the hydrogen atom, making it a hydrogen bond donor
- There must be a hydrogen bond acceptor nearby, which is typically a small, highly electronegative atom (N, O, or F) with at least one lone pair of electrons. The lone pair on this acceptor atom interacts with the partially positive hydrogen, forming the hydrogen bond
- The spatial arrangement is important: the angle between the donor atom, the hydrogen, and the acceptor atom tends to be close to 180°, maximizing the strength of the hydrogen bond
- The size of the atoms involved affects hydrogen bonding. Smaller atoms allow closer approach and stronger interaction, which is why N, O, and F are common participants, while larger atoms like chlorine form weaker hydrogen bonds due to steric hindrance
- Significant electronegativity difference between hydrogen and the atom it is bonded to is necessary to create the partial positive charge on hydrogen. For example, PH3 does not form hydrogen bonds because phosphorus and hydrogen have similar electronegativities, so no strong dipole forms
In summary, for a hydrogen atom in a molecule to participate in hydrogen bonding, it must be bonded to a highly electronegative atom (N, O, or F) creating a polarized bond, and there must be a nearby electronegative atom with a lone pair to act as an acceptor, with appropriate geometry and size considerations to enable the interaction
