Experimental dipole moments

The orientation of the dipole moment, experimentally established by the Stark effect (158), can also be compared to the calculated ones (Fig. 1-5). 

It s relatively easy to measure dipole moments experimentally, and values for some common substances are given in Table 10.1. Once the dipole moment is known, it s then possible to get an idea of the amount of charge separation in a molecule. In chloromethane, for example, the experimentally measured dipole moment is /x = 1.87 D. If we assume that the contributions of the nonpolar C-H bonds are small, then most of the chloromethane dipole moment is due to the C-Cl bond. Since the C-Cl bond distance is 178 pm, we can calculate that the dipole moment of chloromethane would be 1.78 X 4.80 D = 8.54 D if the C-Cl bond were ionic (that is, if a full negative charge on chlorine were separated from a full positive charge on carbon by a distance of 178 pm). But because the measured dipole moment of chloromethane is only 1.87 D, we can conclude that the C-Cl bond is only about (1.87/8.54)(100%) = 22% ionic. Thus, the chlorine atom in chloromethane has an excess of about 0.2 electron, and the carbon atom has a deficiency of about 0.2 electron. 

You could establish die geometiy of HgBr2 by measming its dipole moment. If niercmy(II) bromide were bent, it would have a measmable dipole moment. Experimentally, it has no dipole moment and therefore must be linear. 

The arrows show the shift of electron density from the less electronegative carbon atom to the more electronegative oxygen atom. In each case, the dipole moment of the entire molecule is made up of two bond moments, that is, individual dipole moments in the polar C=0 bonds. The bond moment is a vector quantity, which means that it has both magnitude and direction. The measured dipole moment is equal to the vector sum of the bond moments. The two bond moments in CO2 are equal in magnitude. Because they point in opposite directions in a Unear CO2 molecule, the sum or resultant dipole moment would be zero. On the other hand, if the CO2 molecule were bent, the two bond moments would partially reinforce each other, so that the molecule would have a dipole moment. Experimentally it is found that carbon dioxide has no dipole moment. Therefore, we conclude that the carbon dioxide molecule is linear. The linear nature of carbon dioxide has been confirmed through other experimental measurements. 

Dipole Moments Experimental and Calculated Permanent Dipole of a-Chymotrypsin. 

J. Antoslewlcz and D. Porschke, Biochemistry, 28, 10072 (1989). The Nature of Protein Dipole Moments Experimental and Calculated Permanent Dipole of a-Chymotrypsin. 

McClellan, A. L., Tables of Experimental Dipole Moments, W. H. Freeman, San Francisco (1963-74). 

There are higher multipole polarizabilities tiiat describe higher-order multipole moments induced by non-imifonn fields. For example, the quadnipole polarizability is a fourth-rank tensor C that characterizes the lowest-order quadnipole moment induced by an applied field gradient. There are also mixed polarizabilities such as the third-rank dipole-quadnipole polarizability tensor A that describes the lowest-order response of the dipole moment to a field gradient and of the quadnipole moment to a dipolar field. All polarizabilities of order higher tlian dipole depend on the choice of origin. Experimental values are basically restricted to the dipole polarizability and hyperpolarizability [21, 24 and 21]. Ab initio calculations are an imponant source of both dipole and higher polarizabilities [20] some recent examples include [26, 22] ... 

McClellan A L 1963 Tables of Experimental Dipole Moments vol 1 (New York Freeman)... 

Hollenstein H, Marquardt R, Quack M and Suhm M A 1994 Dipole moment function and equilibrium structure of methane In an analytical, anharmonic nine-dimenslonal potential surface related to experimental rotational constants and transition moments by quantum Monte Carlo calculations J. Chem. Phys. 101 3588-602... 

It has already been said that the merits of a method for charge calculation can be assessed mainly by its usefulness in modeling experimental data. Charges from the PEOE procedure have been correlated with Cls-ESCA shifts [28], dipole moments [33], and NMR shifts [34], to name but a few. 

TabU 3-5 Dipole moments calculated for formaldehyde using various basis sets at the experimental geometry,... 

TIk experimentally determined dipole moment of a water molecule in the gas phase is 1.85 D. The dipole moment of an individual water molecule calculated with any of thv se simple models is significantly higher for example, the SPC dipole moment is 2.27 D and that for TIP4P is 2.18 D. These values are much closer to the effective dipole moment of liquid water, which is approximately 2.6 D. These models are thus all effective pairwise models. The simple water models are usually parametrised by calculating various pmperties using molecular dynamics or Monte Carlo simulations and then modifying the... 

Essentially all experimentally measured properties can be thought of as arising through the response of the system to some externally applied perturbation or disturbance. In turn, the calculation of such properties can be formulated in terms of the response of the energy E or wavefunction P to a perturbation. For example, molecular dipole moments p are measured, via electric-field deflection, in terms of the change in energy... 

The calculated electronic distribution leads to an evaluation of the dipole moment of thiazole. Some values are collected in Table 1-7 that can be compared to the experimental value of 1.61 D (158). 

Experimental measurements of dipole moments give size but not direction We normally deduce the overall direction by examining the directions of individual bond dipoles With alkenes the basic question concerns the alkyl groups attached to C=C Does an alkyl group donate electrons to or withdraw electrons from a double bond d This question can be approached by comparing the effect of an alkyl group methyl for exam pie with other substituents... 

The molecular dipole moment is perhaps the simplest experimental measure of charge density in a molecule. The accuracy of the overall distribution of electrons in a molecule is hard to quantify, since it involves all of the multipole moments. Experimental measures of accuracy are necessary to evaluate results. The values for the magnitudes of dipole moments from AMI calculations for a small sample of molecules (Table 4) indicate the accuracy you may... 

Energy, geometry, dipole moment, and the electrostatic potential all have a clear relation to experimental values. Calculated atomic charges are a different matter. There are various ways to define atomic charges. HyperChem uses Mulliken atomic charges, which are commonly used in Molecular Orbital theory. These quantities have only an approximate relation to experiment their values are sensitive to the basis set and to the method of calculation. 

Tables 5.17 and 5.18 contain a selected group of compounds for which the dipole moment is given. An extensive collection of dipole moments (approximately 7000 entries) is contained in A. L. McClellan, Tables of Experimental Dipole Moments, W. H. Freeman, San Francisco, 1963. A critical survey of 500 compounds in the gas phase is given by Nelson, Tide, and Maryott, NSRDS-NBS 10, Washington, D.C., 1967.

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