INTERMOLECULAR FORCES AND INTERACTION ENERGY 215 



and here the clots represent the dipolar interaction. If this is the correct 

 mechanism, the bond energy should be given approximately by a calcula- 

 tion of this dipolar interaction. Let us take the 0— H ••• bond as it occurs 

 in water: the bond distance is 2.76 A and the distance between dipole cen- 

 ters is therefore 2.40 A (taking into account the configuration of the water 

 molecules), the dipole moment of the 0— H bond is 1.51 debyes and of the 

 water molecule 1.83 debyes, and the 0— H bond distance is 0.955 A. Using 

 Eq. 6-17 the potential energy is found to be — 6.61 kcal mole; this must 

 be corrected for repulsion interaction (see page 231) by a reduction of ap- 

 proximately 35% and the final calculated energy is therefore — 4.29 kcal/ 

 mole. The experimental bond energy is — 4.5 kcal/mole and the agreement 

 is satisfactory; there are, of course, other interactions of less importance, 

 such as induction effects, which will increase the stability slightly. One 

 can therefore interpret the hydrogen bond simply as an ordinary dipole- 

 dipole interaction. Such bonds are formed more readily with hydrogen than 

 with other atoms because the small size of the hydrogen atom allows 

 close approach of the interacting electronegative atoms. 



An interesting discussion of the various theories of the hydrogen bond 

 has been presented by Pimentel and McClellan (1960, p. 226). A critical 

 evaluation of the electrostatic theory was given and it was concluded that 

 it does not account for all the phenomena observed in hydrogen bond for- 

 mation. They believe that a molecular orbital treatment, involving both 

 bonding and nonbonding orbitals, might be preferable, but at the present 

 time such a description cannot be quantitative. Another approach is to 

 interjDret the stability of the hydrogen bond as the result of resonance 

 between structures such as the following: 



where the daslies indicate covalent bonds (Pauling. 1960, p. 449; Wheland, 

 1955, p. 47). These and other theories are usually not as incompatible as 

 has often been assumed. 



The strength of the hydrogen bond will depend in part on the magnitude 

 of the dipole moments involved and thus on the degree of electronegativity 

 of the atoms bonded. The electronegativities for biologically important 

 atoms are (Gordy, 1946): F. 3.95; 0. 3.45; N, 2.98; CI. 2.97; Br, 2,75; C, 

 2.55; S, 2.53; I, 2.45; H, 2.13; and P. 2.10. In Table 6-1 the bond moments 

 and electronegativity differences are given for some X— H bonds of possible 

 significance in inhibitor interactions. A reasonable correlation may be 

 observed. It should be pointed out that these bond moments are average 

 values only and that the actual moment in a molecule will depend on the 



