Hydrogen Bonding: Water
It would seem that the London forces and dipole forces discussed in the earlier sections of this chapter should be adequate to account for macroscopic properties of covalently bonded substances. Certainly they can be successfully applied to hydrocarbons and many polarDescribes a molecule that has separated, equal positive and negative charges that consitute a positive and a negative pole; such a molecule tends to assume certain orientations more than others in an electric field. substances. There are some experimental data, however, which cannot be explained by London and dipoleIn an electrically neutral species, separated, equal positive and negative charges that consitute a positive and a negative pole; such a species tends to assume certain orientations more than others in an electric field. forces alone. An example appears below, where boiling points are plotted for hydrogen compounds (hydrides) of most of the nonmetals.
Hydrides of elements in the fifth period behave as we might predict. SnH4, which consists of nonpolarDescribes a molecule with no net permanent dipole; this can occur when there is no separation of centers of positive and negative electrical charge or because there are bond dipoles that cancel each others' effects. A polar molecule will assume certain orientations more than others in an electric field. molecules, boils at the lowest temperatureA physical property that indicates whether one object can transfer thermal energy to another object.. SbH3, H2Te and HI, all of which are polar, have somewhat higher boiling points, but all lie within a range of 50°C. Similar behavior occurs among the hydrides of elements in the fourth and third periods. In the second period, however, the polar hydrides NH3, H2O, and HF all have boiling points more than 100°C above that of the nonpolar compound CH4. Clearly these second-row hydrides must have particularly strong intermolecular forces.
In order to see why this happens, let us consider the simplest second-row hydride—HF. Suppose that two HF molecules approach each other, as shown in the following figure. In each HF molecule the hydrogen nucleusThe collection of protons and neutrons at the center of an atom that contains nearly all of the atoms's mass. is rather poorly shielded by a thin electronA negatively charged, sub-atomic particle with charge of 1.602 x 10-19 coulombs and mass of9.109 x 1023 kilograms; electrons have both wave and particle properties; electrons occupy most of the volume of an atom but represent only a tiny fraction of an atom's mass. cloud (only two electrons), and much of that electron cloud has been distorted toward the highly electronegative fluorine atom.
The close approach of oppositely charged ends of molecular dipoles combines with this small degree of covalent-bond character to produce an abnormally strong intermolecular force called a hydrogen bondAn attractive force, either intramolecular or intermolecular, between an electronegative atom and a hydrogen atom attached to another electronegative atom.. In order for hydrogen bonding to occur, there must be a hydrogen atom connected to a small, highly electronegative atom (usually fluorine, oxygen, or nitrogen) in one molecule. The other molecule must have a very electronegative atom (again usually fluorine, oxygen, or nitrogen) which has one or more lone pairs of electrons. Separation of two molecules joined by a hydrogen bond requires 10 to 30 kJ mol–1, roughly 10 times the energyA system's capacity to do work. needed to overcome dipole forces. Thus hydrogen bonding can account for the unusually high boiling points of NH3, H2O, and HF.
Hydrogen bonding between HF molecules is particularly evident in solidA state of matter having a specific shape and volume and in which the particles do not readily change their relative positions. HF, where the atomsThe smallest particle of an element that can be involved in chemical combination with another element; an atom consists of protons and neutrons in a tiny, very dense nucleus, surrounded by electrons, which occupy most of its volume. are arranged in a zigzag pattern:
Here the distance between hydrogen and fluorine nuclei in different molecules is only 150 pm. If we add the van der Waals radii for hydrogen and fluorine, we obtain an expected hydrogen to fluorine distance of (120 + 135) pm = 255 pm, over 100 pm larger than that observed. Obviously the hydrogen and fluorine atoms in adjacent molecules are not just “touching,” but must be associated in a much more intimate way.
Provisional Redefinition of the Hydrogen Bond
A task force of the International Union of Pure and Applied Chemistry (IUPAC) is proposing a redefinition of the hydrogen bond:
"The hydrogen bond is an attractive interaction between a hydrogen atom from a molecule or a molecular fragment X-H in which X is more electronegative than H, and an atom or a groupThose elements that comprise a single column of the periodic table. Also called family. of atoms in the same or a different molecule, in which there is evidence of bond formation." 
This proposed redefinition would include many more situations where hydrogen bonding appears to be important, including cases where the hydrogen atom is attracted to atoms other than F,O, and N. For example, there appears to be an attraction between the d subshell electrons in the central platinum atom in the complexA central metal and the ligands surrounding it; also called coordination complex. shown in the Figure below, and the hydrogen atoms of an adjacent H2O or the NH3 group on an adjacent platinum complex.
The new formulation recognizes that the hydrogen bond may have some covalent character. It recognizes that hydrogen bonds may be significant in H2S at higher pressures and low temperatures, or that a "dihydrogen bond" may form (where metal hydrides like LiH are the H bond acceptors).