The International Union of Pure and Applied Chemistry Task Group defined hydrogen bonding as an interaction involving charge transfer and orbital interactions rather than merely electrostatic attraction. This evidence-based definition specifies that a donor atom like nitrogen, oxygen, or fluorine holds the protic hydrogen while an acceptor atom provides a lone pair of electrons.
T. S. Moore and T. F. Winmill mentioned the hydrogen bond concept for the first time in 1912 to explain why trimethylammonium hydroxide acts as a weaker base than tetramethylammonium hydroxide. Linus Pauling later credited these scientists with this initial description before Latimer and Rodebush described hydrogen bonding in water in 1920.
Each water molecule participates in an average of 3.59 hydrogen bonds at 25 degrees Celsius according to liquid water simulations. The number decreases to 3.24 at 100 degrees Celsius due to increased molecular motion and rises to 3.69 at 0 degrees Celsius.
The distance between atoms in water typically measures around 197 picometers when forming hydrogen bonds. Crystallography remains the most important method for identifying these bonds in complicated molecules by showing structural details like distances smaller than van der Waals radii.
Ice forms an open hexagonal lattice stabilized by hydrogen bonds which results in a density less than that of liquid water at the same temperature. This unique property allows the solid phase to float on the liquid because oxygen atoms possess two lone pairs and two hydrogen atoms allowing a total of four bonds per molecule.