Polar molecules electron distribution

Applied electric fields, whether static or oscillating, distort (polarize) the electron distribution and nuclear positions in molecules. Much of this volume describes effects that arise from the electronic polarization. Nuclear contributions to the overall polarization can be quite large, but occur on a slower time-scale than the electronic polarization. Electronic motion can be sufficiently rapid to follow the typical electric fields associated with incident UV to near IR radiation. This is the case if the field is sufficiently off resonance relative to electronic transitions and the nuclei are fixed (see ref 5 for contributions arising from nuclear motion). Relaxation between states need not be rapid, so... 

The electric field of ions can induce dipole moments in uncharged, nonpolar molecules by distorting (polarizing) the electron distribution within the molecules. 

Trichloromethane, CHCI3, is a polar molecule. The three C—Cl dipoles point in a similar direction. Their combined effect is not cancelled out by the polarity of the C—H bond. This is because the C—H bond is virtually non-polar. The electron distribution is asymmetric. The molecule is polar, with the negative end towards the chlorine atoms. This is shown in Figure 4.33a. 

It is advantageous if the laser system pemiits rotation of the optical polarization. Detached electrons correlated witii different final electronic states of the neutral molecule will generally be emitted with different angular distributions about the direction of polarization. Measurement of the angular distribution helps in the interpretation of complex photoelectron spectra. The angular distribution/(0) of photoelectrons is [50]... 

Induced dipole/induced dipole forces are the only intermolecular attractive forces available to nonpolar molecules such as alkanes In addition to these forces polar molecules engage m dipole-dipole and dipole/mduced dipole attractions The dipole-dipole attractive force is easiest to visualize and is illustrated m Figure 4 3 Two molecules of a polar substance experience a mutual attraction between the positively polarized region of one molecule and the negatively polarized region of the other As its name implies the dipole/induced dipole force combines features of both the induced dipole/mduced dipole and dipole-dipole attractive forces A polar region of one mole cule alters the electron distribution m a nonpolar region of another m a direction that produces an attractive force between them... 

Many other measures of solvent polarity have been developed. One of the most useful is based on shifts in the absorption spectrum of a reference dye. The positions of absorption bands are, in general, sensitive to solvent polarity because the electronic distribution, and therefore the polarity, of the excited state is different from that of the ground state. The shift in the absorption maximum reflects the effect of solvent on the energy gap between the ground-state and excited-state molecules. An empirical solvent polarity measure called y(30) is based on this concept. Some values of this measure for common solvents are given in Table 4.12 along with the dielectric constants for the solvents. It can be seen that there is a rather different order of polarity given by these two quantities. 

One area where the concept of atomic charges is deeply rooted is force field methods (Chapter 2). A significant part of the non-bonded interaction between polar molecules is described in terms of electrostatic interactions between fragments having an internal asymmetry in the electron distribution. The fundamental interaction is between the Electrostatic Potential (ESP) generated by one molecule (or fraction of) and the charged particles of another. The electrostatic potential at position r is given as a sum of contributions from the nuclei and the electronic wave function. 

Most organic compounds are electrically neutral they have no net charge, either positive or negative. We saw in Section 2.1, however, that certain bonds within a molecule, particularly the bonds in functional groups, are polar. Bond polarity is a consequence of an unsymmetrical electron distribution in a bond and is due to the difference in electronegativity of the bonded atoms. 

In the HF molecule, the distribution of the bonding electrons is somewhat different from that found in H2 or F2. Here the density of the electron doud is greater about the fluorine atom. The bonding electrons, on the average, are shifted toward fluorine and away from the hydrogen (atom Y in Figure 7.9). Bonds in which the electron density is unsymmetrical are referred to as polar bonds. 

In Section 2.12, we saw that a polar covalent bond in which electrons are not evenly distributed has a nonzero dipole moment. A polar molecule is a molecule with a nonzero dipole moment. All diatomic molecules are polar if their bonds are polar. An HC1 molecule, with its polar covalent bond (8+H—Clfi ), is a polar molecule. Its dipole moment of 1.1 D is typical of polar diatomic molecules (Table 3.1). All diatomic molecules that are composed of atoms of different elements are at least slightly polar. A nonpolar molecule is a molecule that has no electric dipole moment. All homonuclear diatomic molecules, diatomic molecules containing atoms of only one element, such as 02, N2, and Cl2, are nonpolar, because their bonds are nonpolar. 

A satisfactory feature of (116) is that the first-order term, which describes the interaction between the external charge and the unperturbed molecule, continues to disappear if the electron distribution is highly uniform as in an alternant hydrocarbon in such cases the propensities for reaction are still dominated by the polarization term, but this now has a less simple form, depending upon all atom-atom polarizabilities and on the position of the attacking ion with respect to all conjugated centres. 

Two atoms of the same electronegativity will share electrons equally in a pure covalent bond therefore, any molecule that contains atoms of only one element, like H2 or CI2, has pure covalent bonding. Two atoms of different electronegativities, however, will have either the distorted electron distribution of a polar bond or the complete electron transfer of an ionic bond. Table 5-6 interprets the bonding between two elements as a function of the difference in their electronegativity. 

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