Diatomic molecules dipole moments

In this chapter we will also focus on the dipole moment of molecules. With these, some of the most interesting phenomena are the molecules for which the electric moment is in the wrong direction insofar as the atomic electronegativities are concerned. CO is probably the most famous of these cases, but other molecules have even more striking disagreements. One of the larger is the simple diatomic BF. We will take up the question of the dipole moments of molecules like BF in Chapter 12. In this chapter we will examine in a more general way how various sorts of structures influence electric moments for two simple cases. For some of the discussion in this chapter we restrict ourselves to descriptions of minimal basis set results, since these satisfactorily describe the physics of the effects. In other cases a more extensive treatment is necessary. 

Experimental data has been taken from the following compendia R.D. Nelson, D.R. Tide and A.A. Maryott, Selected Values of Electric Dipole Moments for Molecules in the Gas Phase, NSRDA-NBS 10, U.S. Government Printing Office, Washington, D.C., 1967 (b) A.L. McClellan, Tables of Experimental Dipole Moments, W.H. Freeman, San Francisco, 1963 (c) ibid., vol. 2, Rahara Enterprises, El Cerritos, CA, 1974 (d) K.P. Huber and G Herzberg, Molecular Spectra and Molecular Structure. IV. Constants of Diatomic Molecules, Van Nostrand Reinhold, New York, 1979. 

Of course, any diatomic (two-atom) molecule that has a polar bond also will show a molecular dipole moment Polyatomic molecules also can exhibit dipolar behavior. For example, because the oxygen atom in the water molecule has a greater electronegativity... 

Although a diatomic molecule can produce only one vibration, this number increases with the number of atoms making up the molecule. For a molecule of N atoms, 3N-6 vibrations are possible. That corresponds to 3N degrees of freedom from which are subtracted 3 translational movements and 3 rotational movements for the overall molecule for which the energy is not quantified and corresponds to thermal energy. In reality, this number is most often reduced because of symmetry. Additionally, for a vibration to be active in the infrared, it must be accompanied by a variation in the molecule s dipole moment. 

If one of the components of this electronic transition moment is non-zero, the electronic transition is said to be allowed if all components are zero it is said to be forbidden. In the case of diatomic molecules, if the transition is forbidden it is usually not observed unless as a very weak band occurring by magnetic dipole or electric quadnipole interactions. In polyatomic molecules forbidden electronic transitions are still often observed, but they are usually weak in comparison with allowed transitions. 

Figure Bl.2.1. Schematic representation of the dependence of the dipole moment on the vibrational coordinate for a heteronuclear diatomic molecule. It can couple with electromagnetic radiation of the same frequency as the vibration, but at other frequencies the interaction will average to zero.

Laser Raman diagnostic teclmiques offer remote, nonintnisive, nonperturbing measurements with high spatial and temporal resolution [158], This is particularly advantageous in the area of combustion chemistry. Physical probes for temperature and concentration measurements can be debatable in many combustion systems, such as furnaces, internal combustors etc., since they may disturb the medium or, even worse, not withstand the hostile enviromnents [159]. Laser Raman techniques are employed since two of the dominant molecules associated with air-fed combustion are O2 and N2. Flomonuclear diatomic molecules unable to have a nuclear coordinate-dependent dipole moment caimot be diagnosed by infrared spectroscopy. Other combustion species include CFl, CO2, FI2O and FI2 [160]. These molecules are probed by Raman spectroscopy to detenuine the temperature profile and species concentration m various combustion processes. 

This is the same as Equation (5.14) for a diatomic or linear polyatomic molecule and, again, the transitions show an equal spacing of 2B. The requirement that the molecule must have a permanent dipole moment applies to symmetric rotors also. 

Equation (6.8), to (d /dx)g. Figure 6.1 shows how the magnitude /r of the dipole moment varies with intemuclear distance in a typical heteronuclear diatomic molecule. Obviously, /r 0 when r 0 and the nuclei coalesce. For neutral diatomics, /r 0 when r qg because the molecule dissociates into neutral atoms. Therefore, between r = 0 and r = oo there must be a maximum value of /r. Figure 6.1 has been drawn with this maximum at r < Tg, giving a negative slope d/r/dr at r. If the maximum were at r > Tg there would be a positive slope at r. It is possible that the maximum is at r, in which case d/r/dr = 0 at Tg and the Av = transitions, although allowed, would have zero intensity. 

Figure 6.1 Variation of dipole moment fi with intemuclear distance r in a heteronuclear diatomic molecule...

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 diatomic molecule is polar if its bond is polar. A polyatomic molecule is polar if it has polar bonds arranged in space in such a way that the dipole moments associated with the bonds do not cancel. 

Dipole moments also depend on molecular shape. Any diatomic molecule with different atoms has a dipole moment. For more complex molecules, we must evaluate dipole moments using both bond polarity and molecular shape. A molecule with polar bonds has no dipole moment if a symmetrical shape causes polar bonds to cancel one another. 

Let us suppose that the system of interest does not possess a dipole moment as in the case of a homonuclear diatomic molecule. In this case, the leading term in the electric field-molecule interaction involves the polarizability, a, and the Hamiltonian is of the form ... 

As mentioned earlier, heavy polar diatomic molecules, such as BaF, YbF, T1F, and PbO, are the prime experimental probes for the search of the violation of space inversion symmetry (P) and time reversal invariance (T). The experimental detection of these effects has important consequences [37, 38] for the theory of fundamental interactions or for physics beyond the standard model [39, 40]. For instance, a series of experiments on T1F [41] have already been reported, which provide the tightest limit available on the tensor coupling constant Cj, proton electric dipole moment (EDM) dp, and so on. Experiments on the YbF and BaF molecules are also of fundamental significance for the study of symmetry violation in nature, as these experiments have the potential to detect effects due to the electron EDM de. Accurate theoretical calculations are also absolutely necessary to interpret these ongoing (and perhaps forthcoming) experimental outcomes. For example, knowledge of the effective electric field E (characterized by Wd) on the unpaired electron is required to link the experimentally determined P,T-odd frequency shift with the electron s EDM de in the ground (X2X /2) state of YbF and BaF. 

Values of the dipole moment of some diatomic molecules are given in Table 2.11. The SI unit of dipole moment is the coulomb-meter (C-m). This is a very large unit, so in Table 2.11 we use the unit 10-30 C-m. Dipole moments are often quoted in an older unit, the debye (D) 1 D = 3.24 X 1CT30 C-m. We can see from Table 2.11 that the dipole moment of a diatomic molecule usually reflects the difference between the electronegativities of the two atoms. 

Because the dipole moment of a diatomic molecule is qd, it would appear that if we knew the interatomic distance (bond length) d, we should be able to calculate the atomic charges q. For example, the bond length of the HC1 molecule is 127 pm and the dipole moment is 3.44 X 10 3n C-m, so we have... 

It has often been assumed that atomic charges can be calculated from the measured dipole moment of a diatomic molecule and the bond length. For this assumption to hold, however, the center of negative charge of an atom would have to be situated at the nucleus, in other words, atoms would have to be spherical. But we have seen that atoms in molecules are not spherical, and so the center of negative charge is not centered at the nucleus. Each atom therefore has a dipole moment called the atomic dipole moment (Chapter 2). 

As mentioned above and discussed in Chapter 2, atomic charges were often obtained in the past from dipole moments of diatomic molecules, assuming that the measured dipole moment equal to the bond length times the atomic charge. This method assumes that the molecular electron density is composed of spherically symmetric electron density distributions, each centered on its own nucleus. That is, the dipole moment is assumed to be due only to the charge transfer moment Mct. and the atomic dipoles Malom are ignored. 

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