Electronegativity and Molecular Dipoles

<------ higher positive charge

higher negative charge ----->

Linus Pauling established a scale of electronegativities for the atoms. The electronegativities common in organic chemistry are displayed on the right as they are in the Periodic Table. The greater the number, the greater the electronegativity. Elements in the first group (vertical) are highly electropositive. The lighter the shade of blue in Group I, the more electropositive. This trend is true for Groups II-VII as well. As one moves across the periods of the table (horizontal), the atoms become more electronegative. Ionic compounds are formed from highly electropositive and electronegative atoms, e. g., Na⁺Cl^-. Since carbon is in the center of the second period, it forms covalent bonds with its near neighbors. Thus, a covalent C-H bond is often displayed as (δ^-) C-H (δ⁺), indicating that there is partial negative charge on carbon and partial positive charge on hydrogen. Similarly, a carbon-boron bond would be polarized (δ^-) C-B (δ⁺), However, a boron-hydrogen bond would be polarized in the opposite sense, (δ⁺) B-H (δ^-).

Bond Dipole Moments:

The bond dipole moment (BDM) is defined as μ = q x d, where q is the partial charge on either atom times the distance between the charges. In the cgs system charge is measured in statcoulombs (esu, electrostatic units) and distance in centimeters. A BDM of 10^-10 statcoulombs times 10^-8 cm = 10^-18 statcoulomb-centimeters. This quantity is defined as 1 Debye (1D), named in honor of a pioneer in this field, Peter Debye. Also note, 10^-8 cm = 1 Ångstrom (Å). The charge on a proton or an electron is 4.8 x 10^-10 statcoulomb.

Pauling Electronegativities

If these two charges were held 1Å apart, the BDM would be 4.8D. BDM's may also be expressed in the MKS system where 1 coulomb = 3 x 10⁹ statcoulombs and distance is measured in meters. The BDM is designated by the arrow shown in the Table on the right, with the arrow pointing in the direction of the negative charge. All BDM's are listed with the more electronegative atom to the right.

Molecular Dipole Moments:

Molecular dipole moments (MDM) are measurable quantities. The BDM's in the previous section are derived from the MDM's knowing the geometry of the molecule. In other words, the molecular dipole is the vector sum of the BDM's. Symmetrical molecules like H₂, CO₂, methane and CCl₄ have no net molecular dipole moment.

Electrostatic Potential Maps:

These maps are a measure of regions of low and high electric potential in a molecule. They provide an insight into the location of electron density. Regions of high negative potential are red; low potential is in blue. Intermediate regions follow the color code in the Electronegativity Table.

Bond Dipole Moments

Bond	Dipole Moment (D)
H-C	0.3
H-N	1.31
H-O	1.53
C-N	0.22
C-O	0.86
C-F	1.51
C-Cl	1.56
C-Br	1.48
C-I	1.29
C=O	2.4
CN (cyano)	3.6

Ethane, Ethylene, Acetylene:

These three two-carbon molecules have no molecular dipole. They have sp³ (tetrahedral), sp² (planar) and sp (linear) hybridized carbons, respectively. What is the conformation of ethane? In ethylene, the highest electron density (red) is in the center of the molecule, above and below the plane of the molecule. This is due to the π-bond. In acetylene, the red region extends about the "equator" of the molecule, owing to the orthogonal pair of π-bonds.

The four methyl halides shown above have increasing (left to right) electronegative halogen atoms. View the halogen atoms along the X-C axis. Notice the region of the halogen atom along the bond axis but remote from the bond. There is less and less electron density at this site as one passes from fluorine to iodine.
The green region of the chart on the right gives the experimentally obtained quantities: dipole moment and bond length. The calculated dipole molecular moment is the product of the bond length times the charge on an isolated electron, 4.8 x 10^-10 statcoulomb. Clearly, the calculated values are too high. The charge on the carbon and halide must be a fractional charge. Dividing the experimental dipole by the calculated value provides the fractional electron charge.