Ammonia molecular dipole moment

Another fundamental distinction between H and Li bonds is associated with the charge distributions occasioned by formation of the complex. Szczesniak et al. [191] explored these redistributions by way of spectroscopic atomic charges which act to mimic experimental quantities such as vibrational intensities. They noted that whereas most of the charge extracted from the ammonia in H3N HCl was picked up by the Cl atom, it is the Li atom in H3N LiCl that is the ultimate sink of electron density. The authors were able to discern a relationship also between the total charge transferred from NH3 to the electron acceptor and the calculated intensity of the intermolecular stretch. This finding conforms to the requirement of a changing molecular dipole moment in order to lend intensity to this vibration. 

The presence of a lone pair has a significant effect on the molecular dipole moment. The two electrons of a lone pair are balanced by two positive charges in the nucleus, but the lone pair is separated from the nucleus by some distance. There is, therefore, a dipole moment associated with every lone pair. Common examples are ammonia and water (Figure 1.46). 

In fact, in the ammonia synthesis reaction, there is an additional promoting effect due to the addition of an alkali (see Fig. 7.7). Intermediate NH species formed during the ammonia synthesis reactim are destabrilized at the surface. The H atoms tend to be slightly positive, and hence, the molecular dipole moment has the opposite sign. This gives... 

As we saw in Chapter 2, ammonia has a pyramidal structure with a lone pair of electrons that is associated with its basicity (4.3). The arrow next to the molecule shows the direction of the resultant molecular dipole moment. One of the earliest pieces of evidence for the structure of ammonia was that it has a dipole moment, symbol p. The N-H bond is polarized such that there is a small positive charge on hydrogen and a small negative charge on nitrogen. The resultant molecular dipole when these three bond dipoles are summed lies along the threefold symmetry axis of the molecule. For those of you who have studied vector algebra, this is a vector sum. 

Physical properties of the solvent are used to describe polarity scales. These include both bulk properties, such as dielectric constant (relative permittivity), refractive index, latent heat of fusion, and vaporization, and molecular properties, such as dipole moment. A second set of polarity assessments has used measures of the chemical interactions between solvents and convenient reference solutes (see table 3.2). Polarity is a subjective phenomenon. (To a synthetic organic chemist, dichloromethane may be a polar solvent, whereas to an inorganic chemist, who is used to water, liquid ammonia, and concentrated sulfuric acid, dichloromethane has low polarity.)... 

Not surprisingly, the largest dipole moment listed in Table 10.1 belongs to the ionic compound NaCl. Water and ammonia also have substantial dipole moments because both oxygen and nitrogen are electronegative relative to hydrogen and because both O and N have lone pairs of electrons that make substantial contributions to net molecular polarity ... 

Water is also included in the table to make one point— the solvent that we are all most familiar with is a poor candidate from both engineering and safety standpoint. The critical temperature and pressure are among the highest for common solvents. Ammonia is very unpleasant to work with since a fume hood or other venting precautions are needed to keep it out of the laboratory atmosphere. One of the alternative fluids of potential interest is nitrous oxide. It is attractive since it has molecular weight and critical parameters similar to carbon dioxide, yet has a permanent dipole moment and is a better solvent than carbon dioxide for many solutes. There are evidences of violent explosive reactions of nitrous oxide in contact with oils and fats. For this reason, nitrous oxide should be used with great care and is not suitable as a general purpose extraction fluid. 

In the first two parts of this chapter, electron transfer (ET) from atomic donors, e.g., alkali metals or the iodine anion, to an accepting unit composed of simple molecular or atomic solvents was discussed. It was demonstrated that even for a molecule without a stable anionic state or large dipole moment, e.g., water and ammonia, an ensemble of a relatively small number of the molecules can act as an electron acceptor. In the case of the solvated alkali metal atom clusters, ET takes place spontaneously as the number of solvent molecules increases, while the ET in the solvated 1 clusters is induced by photoexcitation into the diffuse electronic excited states just below the vertical detachment thresholds. These ET processes in isolated supermolecular systems resemble the charge delocalization phenomena in condensed phases, e.g., excess-electron ejection from alkali metals into polar solvents and the charge transfer to solvent in a solution of stable anions. 

It is also worthwhile to compare the 6-31G values with those that would be obtained if perfect molecular orbitals were used. Neumann and Moskowitz33,34 calculated dipole moments for water and formaldehyde of 1.995 and 2.83 debyes, respectively. More recently, Dunning, Pitzer, and Aung35 have performed a similar study for water and obtained 2.03 debyes. Near Hartree-Fock wave functions have also been reported for ammonia by Laws, Stevens, and Lipscomb.36 Their best wave function gives a calculated dipole moment of 1.687 debyes. These limited data suggested that the 6-31G results are probably within 0.1-0.2 debye of the Hartree-Fock limit. 

NH3 dipole moment in the molecular bond representation. The first global description of the dipole moments of ammonia was reported by Marquardt et al. [66] who computed ab initio DMSs of NH3 using the MP2/ aug-cc-pVQZ and MCSCF/aug-cc-pVQZ levels of theory. This was also the first work where the MB-representation was used for describing DMSs of an XY3 molecule, formulated as follows ... 

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