ABSTRACT
One can see that even in this official definition the unique property of hydrogen is underlined, namely, its small size that in turn provides the proximity of interacting dipoles. The energy for the typical hydrogen bonds ranges between 0.13-0.31 eV [26]. A classical example is the formation of hydrogen bonds between two water molecules (see Fig. 2.1), which results in the formation of a water dimer (H2O)2 with bond energy of 0.2 eV [26]. The hydrogen bond is depicted by the stroke line. If there are two water molecules one of the hydrogen would tend to interact with both oxygens (Fig. 2.1); as a result a new compound is formed where the dots denote a new type of intermolecular interaction namely hydrogen bond. The hydrogen bond is related to weak bonds, since its energy contains approximately one twentieth of the covalent O-H bond energy. The distance between two oxygen atoms in the complex is approximately equal to 2.76 Å which is less than the sum of the van der Waals radii of oxygen atoms equal to 3.06 Å [26]. The van der Waals radius, rw, of an atom is the radius of an imaginary hard sphere which can be used to model the atom for many purposes.
Figure 2.1 Formation of a water dimer via the hydrogen bond. Since the covalent bond of the hydrogen atom is saturated, the lone electron pair is donated by the atom Y to form the hydrogen bond with it; thus the hydrogen bond being a case of donor-acceptor or coordination bond. As it was mentioned in Chapter 1, a hydrogen atom is always inclined to play the role of a Lewis acid, thus of an acceptor of electron density. For this reason the hydrogen bond is more intensive if the hydrogen atom is bonded by the covalent bond to another atom Y with high electronegativity (e.g., oxygen or fluorine); then the hydrogen atom acquires partial positive charge that enhances its Lewis acidity. Due to this fact the electron density on the hydrogen
