Dipole ionic molecules

Figure 2.4 The dipole moment of a hypothetical purely ionic molecule with spherical ions.

Figure 2.5 (a) Hypothetical ionic molecule with spherical ions and the corresponding charge transfer moment, (b) In a real molecule the ions are polarized leading atomic dipoles in each atom that oppose the charge transfer moment. 

Forty-four five-membered heterocycles of type A (13, 19) have been described (Table I). If the atoms or groups a, b, c, d, e, and f are selected from suitably substituted carbon, nitrogen, oxygen, and sulfur atoms, then with these conditions it can be shown that 144 structural possibilities are provided by the general formula 19. The number of structural possibilities can be deduced in various ways, but a very useful approach is to regard type A meso-ionic molecules (19) as being derived by the union (-<—u— ) of 1,3-dipoles (34) and heterocumulenes (35). 

The STO-3G model provides a very non-uniform account of dipole moments in these compounds (see Figure 10-1). Calculated dipole moments for extremely polar ( ionic ) molecules like lithium chloride are almost always much smaller than experimental values, while dipole moments for moderately polar molecules such as silyl chloride are often larger, and dipole moments for other molecules like carbon... 

All density functional models exhibit similar behavior with regard to dipole moments in diatomic and small polyatomic molecules. Figures 10-6 (EDFl) and 10-8 (B3LYP) show clearly that, except for highly polar (ionic) molecules, limiting (6-311+G basis set) dipole moments are usually (but not always) larger than experimental values. 

Individual errors are typically quite small (on the order of a few tenths of a debye at most), and even highly polar and ionic molecules are reasonably well described. Comparison of results from 6-3IG and 6-311+G density functional models (Figure 10-5 vs. 10-6 for the EDFl model and Figure 10-7 vs. 10-8 for the B3LYP model) clearly reveals that the smaller basis set is not as effective, in particular with regard to dipole moments in highly polar and ionic molecules. Here, the models underestimate the experimental dipole moments, sometimes by 1 debye or more. 

Because the dispersion force acts between neutral molecules it is ubiquitous (compare the gravitational force) however, between polar molecules there are also other forces. Thus, there may be permanent dipole-dipole and dipole-induced dipole interactions and, of course, between ionic species there is the Coulomb interaction. The total force between polar and non-polar (but not ionic) molecules is called the van der Waals force. Each component can be described by an equation of the form V = C/rf, where for the dipole-dipole case n = 6 and C is a function of the dipole moments. Clearly, it is easy to give a reasonable distance dependence to an interaction however, the real difficulty arises in determining the value of C. 

Electrostatic bonding is much more common than covalent bonding in drug-receptor interactions. Electrostatic bonds vary from relatively strong linkages between permanently charged ionic molecules to weaker hydrogen bonds and very weak induced dipole interactions such as van der Waals forces and similar phenomena. Electrostatic bonds are weaker than covalent bonds. 

The second formula means merely that the HC1 molecule is a resonance hybrid between the ionic molecule H+Cl" and the molecule with the purely covalent bond, the direction of the arrow giving the direction in which the electrons have, on the average, been displaced (66). As, however, such an arrow is used by others (57), for indicating a coordinate link (semipolar double bond) caused by a lone electron pair of the donor atom, which likewise produces a dipole with its positive end on the donor side and its negative one on the acceptor side, the author suggests that the symbol — be used for the normal covalent bond, which, by resonance with an ionic structure, possesses a dipole. The point of this half arrow also indicates the direction of the negative end of the dipole. The full arrow — will then be reserved for the coordinate link. Both links play their roles in chemisorption, and it may be useful for the purposes of this article to introduce relatively simple symbols. According to this principle HC1 should be formulated as H—1-Cl. 

The adsorption of polar molecules on surfaces of ionic crystals (Sec. V,5) is influenced by active spots of the same kind as influence the action of Coulomb forces. The effect of these active spots is, quantitatively, less for dipole-containing molecules than for ions. The effect of dipoles on metal surfaces is small (Sec. V,5), and active spots are not expected to give appreciably higher contributions. 

The ion-dipole interaction energy is usually large so that the ion and the dipole align each other. The dissolution of ionic molecules in polar solvents (e.g., water) is due to the ion-dipole interaction (e.g., ZnCl in polyalcohol). When a dipole attracts toward an oppositely charged ion, the comparable ion-ion bond energy in the solid state is released. 

The vibrational dependence of the molecular constants is summarised by English and Zorn [51] for v = 0, 1, 2 and the results are listed in table 8.12. The electric dipole moment of the CsF molecule is large (over 7 D) but, according to Hughes [48], is not as large as one would expect for a purely ionic molecule. The decrease in the electric quadrupole constant as v increases is attributed by English and Zorn [51] to an increasing asymmetry of the internuclear potential. The measured C3 constant is actually the sum of two separate contributions... 

These agree rather well with the experimental values listed above, suggesting that the ionic model is a good one. On the other hand, the negative Fermi contact constant can only arise through polarisation of the electron spins in a covalent bond between the two atoms. The electric dipole moment also seems to be inconsistent with a purely ionic model, yet the quadrupole coupling constant eq0Q is very close to that of the ionic molecule LiF. 

Because a strong electric field is required to align the molecules, further restrictions are imposed on the molecules they should have a permanent dipole moment. For instance, EFISHG can not be applied to measure the second-order nonlinear susceptibilities of octopolar molecules, even though at the molecular level, then-molecular hyperpolarizability, is non-zero. Also, EFISHG can not be used with ionic molecules or with a polar solvent. 

Problem 5.5 The gaseous HF molecule has a dipole moment of 1.83 Debye. The dissociation energy Do = 566 kJ mopi, the bond distance is 92 pm and the vibrational wavenumber 4138 cm . Use the spherical ion model to estimate the dipole moment and the dissociation energy (in kJ moP ) at 0 K. Would you describe HF as an ionic molecule ... 

The dipole moment of ethane is zero by symmetry. Amine borane, on the other hand, is a very polar molecule with an electric dipole moment of 5.22 Debye, nearly as large as that of the ionic molecule LiF, 6.28 D. 

In lower salinity solvents, particularly in distilled water, the electrostatic field in the immediate vicinity of the ionic groups on the polymer chains is only partially neutralized. A strong electrostatic repulsion exists not only between the ionic groups inside an individual molecule, but between the different polymer polyions. This strong electrostatic repulsion works against the buildup of a polymolecular layer. Also, because of the large number of electrostatically held dipole water molecules, a molecule in an outer adsorption zone would be easier dragged away. 

The Inner Helmholtz Plane (IHP) is an ionic layer that consists of adsorbed dipole H2O molecules. The majority of the anions do not penetrate this layer, some do as indicated in Figure 2.11. The inner potential on the boundary of this ionic plane is x. 

FIGURE 18.5 Peter J. W. Debye (1884-1966) was a Dutch-American physical chemist who made important advances in the understanding of ionic solutions and dipoles in molecules. He also formulated an acceptable theory of the thermodynamic properties of crystals at low temperatures. He was awarded the 1936 Nobel Prize in chemistry for his work. 

Equation 4.2 implies that if R is large, the potential may be taken in a first approximation as just the first non-zero term in the expansion, because further terms depend on increasing inverse powers of distance. So for a non-ionic molecule with non-zero dipole the monopole term is zero, while quadrupole and higher multipole terms are much smaller than the dipole term. At large intermolecular distance, therefore, the electric potential of a molecule is to a good approximation described by just the dipole term. For condensed phases, where molecules may come very close to one another, the dipolar approximation is unsatisfactory. 

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