Lone pair moment

NF3 is a colourless, odourless, thermodynamically stable gas (mp —206.8°, bp —129.0°, AG29g — 83.3kJmol ). The molecule is pyramidal with an F-N-F angle of 102.5°, but the dipole moment (0.234 D) is only one-sixth of that of NH3 (1.47 D) presumably because the N-F bond moments act in the opposite direction to that of the lone-pair moment ... 

The dipole moment of phosphabenzene is reinforced by the methyl group shown in (147) and increases it from 1.46 to 1.77 D thus it resembles pyridine, which has the heteroatom at the negative end of the dipole.181 The magnitudes and trends of the dipole moments of the methylphosphines have been investigated by MO studies and the dipoles partitioned into bond moments, bond polarization, and lone-pair moments.182 The reciprocal effects of the double bond and the phosphorus atom in... 

The dipole moment of a compound is a function of the distribution of charge within the molecule, and appears to be a sensitive test for the accuracy of the compound s molecular wave functions. The dipole moment of a molecule can be approximated for a given direction as the sum of two components, /iq, the contribution from net charge densities on the atoms, and for each atom A, ftsp (A), an atomic polarization moment produced by the distortion of the electronic cloud around the atom. The atomic polarization moment results essentially from the mixture of s and p orbitals and, for a heteroatom, includes mv, the lone pair moment. 

There is no major objection to including the lone pair moment in the carbon-halogen bond, which will as a result lie along the nuclear axis, with the halogen as the negative end. 

The interpretation of dipole moments of aromatic compounds is fraught with difficulties because the measured moment is the sum of r-bond, dipole moments of compounds such as pyridones clearly point to enhanced delocalization, quantitative treatment is not easy. Indeed in some cases the method is misleading, as with the borazabenzenes71 where a- and r-bond moments may operate in opposite directions (Section III,D, 10). 

Pauling has estimated that all three are of comparable importance, which is in accord with the low (0.1 d) dipole moment of the molecule, although the very short bond distance (1.13 A) would suggest that (3-VIc) is predominant. The low dipole moment can be perhaps explained on the basis of smaller contributions from (3-VIa) and (3-VIb) if polarity due to lone-pair moments (page 120) is taken into account. 

The important point illustrated by this discussion of HC1 is that the net molecular dipole moment is a vector sum of the bond moment and the lone-pair moments. There is clearly no necessary relationship between molecular dipole moments and bond polarity alone. Naturally, an entirely analogous situation prevails with polar polyatomic molecules, as illustrated by NH3 and NF3. 

Fig. 3-18. Diagram illustrating how the molecular dipole moment,, can be regarded as a vector sum of bond moments and the lone pair moment in NHj and TSF3.

A P-H bond moment Xb = 0.30D, directed in the P H" sense, was derived from theoretically calculated values for the total dipole moment i = -0.66 (see below) and for a lone-pair contribution of -1.15 D (due to the 5ai orbital having mainly P3p character see p. 142). The resultant of the three bond moments (0.49 D) was thus directed oppositely to the total dipole moment [14]. Lone-pair moments of -0.4 or -0.2D (and small bond moments) were derived from the experimental value and its variation with the symmetry coordinate S2 (for two different signs of a x/aS2) [37]. A lone-pair moment of only 0.2 D was used in an investigation of the influence of the inductive effect on the bond moments in a series of phosphanes [16]. Bond moments were also derived assuming a vanishing lone-pair moment (and using an experimental value jx = 0.58) Pb = 0.3574 D [17] and 0.36 D [18, 19]. 

Various sizes of basis sets, especially an inclusion of d functions, and/or various methods of calculation were tested in each of the following papers [14, 25 to 35] see also theoretical work listed in the section dealing with atomic charges, p. 142. For a calculation of bond and lone-pair moments, see also [36]. 

The bond dipole moment i,0F = 0-239 D was obtained from a total measured dipole moment (0.297 D) by assuming ito be the vector sum of the bond moments [25]. If, however, a lone-pair moment at 0 of 3.2 D is assumed (as in the case of HgO), then a true value i,op = 2.9 or 2.3 D results, depending upon the sign of the total moment (taken as ( )0.4 D) [7]. The derivative with respect to the internuclear distance, = D/A, was determined from IR intensities [25]. 

Phosphine is a colourless gas at room temperature, boiling point 183K. with an unpleasant odour it is extremely poisonous. Like ammonia, phosphine has an essentially tetrahedral structure with one position occupied by a lone pair of electrons. Phosphorus, however, is a larger atom than nitrogen and the lone pair of electrons on the phosphorus are much less concentrated in space. Thus phosphine has a very much smaller dipole moment than ammonia. Hence phosphine is not associated (like ammonia) in the liquid state (see data in Table 9.2) and it is only sparingly soluble in water. 

Because of the presence of the lone pairs of electrons, the molecule has a dipole moment (and the liquid a high permittivity or dielectric constant). 

Moleeular properties sueh as dipole moment and polarizability, although in eertain fully empirieal models, bond dipoles and lone-pair eontributions have been ineorporated (although again only for eonventional ehemieal bonding situations). 

In H2O and NH3, shown in Figures 4.18(a) and 4.18(b), the direction of the dipole moment is along the C2 or C3 axis, respectively. In both molecules there are lone pairs of electrons directed away from the 0-FI or N-FI bonds so that the negative end of the dipole is as shown in each case. 

The method has also been applied to partially saturated systems. For instance, the dipole moments of a series of 1-acylpyrazolines (42) with R = H, Me, Et and Ph have been measured (72CHE445) they range from 3.46 to 4.81 D. When compared with values computed by the fragmentary calculation method, the conclusion was reached that here also the E form predominates. In all these examples the lone pair-lone pair repulsions determine the most stable conformation. 

The exact expression for the dipole moment does n( consider atoms as point charges, but rather as nuclei (eat with a positive charge equal to the atomic number) ar electrons (each with unit negative charge). Atoms wii lone pairs may contribute to the dipole moment, even the atom is neutral, as long as the lone pair electrons a not symmetrically placed around the nucleus. 

In contrast with water, methanol, ammonia, and other substances in Table 2.1, carbon dioxide, methane, ethane, and benzene have zero dipole moments. Because of the symmetrical structures of these molecules, the individual bond polarities and lone-pair contributions exactly cancel. 

Methylamine contains an electronegative nitrogen atom with two lone-pair electrons. The dipole moment thus points generally from -CH3 toward NH2. 

If the four atoms attached to the central atom in a tetrahedral molecule are the same, as in tetrachloromethane (carbon tetrachloride), CCI4 (30), the dipole moments cancel and the molecule is nonpolar. However, if one or more of the atoms are replaced by different atoms, as in trichloromethane (chloroform), Cl ICI, or by lone pairs, as in NH3, then the dipole moments associated with the bonds are not all the same, so they do not cancel. Thus, the CHCI, molecule is polar (31). 

Each molecule has four fluorine atoms at the comers of a square. The Xe—E bond polarities cancel in pairs, leaving XeFq with no dipole moment. Four bond polarities also cancel in CIF5, but the fifth Cl—F bond has no counterpart in the opposing direction, so CIF5 has a dipole moment that points along the axis containing the lone pair and the fifth Cl—F bond. 

XeFq, like I3, is a species that has lone pairs but zero dipole moment. Symmetry explains this. We can see why by examining how the lone pairs are placed. In I3, the three lone pairs form a symmetrical trigonal plane. In XeFq, the two lone pairs oppose each other. 

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