Polarity of Bonds and Molecules

Bond Polarity

Polarity in organic chemistry refers to a separation of charge and can describe a bond or an entire molecule. Experimentally, bond polarity is measured by its dipole moment. Bonds connecting atoms of different electronegativity are polar with a higher density of bonding electrons around the more electronegative atom giving it a partial negative charge (designated as d-). The less electronegative atom has some of its electron density taken away giving it a partial positive charge (d+).

This polarization of charge in the H-Cl bond is due to different electronegativities of chlorine and hydrogen.

hcl.gif (3446 bytes)

The polarity in the bond can also be represented by a arrow indicating a dipole (two charges separated by a distance). The tip of the arrow points toward the more electronegative atom.

Molecular Polarity

The polarity of the molecule is the sum of all of the bond polarities in the molecule. Since the dipole moment (m, measured in Debyes (D)) is a vector (a quantitiy with both magnitude and direction), the molecular dipole moment is the vector sum of the individual dipole moments. If we compare the molecular dipole moments of formaldehyde and carbon dioxide, both containing a polar carbonyl (C=O) group, we find that formaldehyde is highly polar while carbon dioxide is nonpolar . Since CO2 is a linear molecule, the dipoles cancel each other.

co-dp.gif (2319 bytes)

Water is a bent molecule with polar O-H bonds. The bond dipole moments add to give a resultant dipole (m = 1.85 D) directed toward the more electronegative oxygen.

The polarity of chloromethanes reveals the importance of symmetry. All of these compounds contain polar C-Cl bonds but the tetrahedral symmetry of CCl4 causes the bond dipoles to cancel giving a nonpolar molecule.

 

ch3cl-bd.gif (4402 bytes) ch2cl2-bd.gif (5622 bytes) chcl3-bd.gif (6492 bytes) ccl4-bd.gif (7709 bytes)
ch3cl-dp.gif (1153 bytes) ch2cl2-dp.gif (1447 bytes) chcl3-dp.gif (1390 bytes) ccl4-dp.gif (1142 bytes)
Chloromethane

The top image show the bond electron density and the bottom image the molecular dipole.

m = 1.87 D

Dichloromethane

The top image show the bond electron density and the bottom image the molecular dipole

m = 1.54 D

Trichloromethane

The top image show the bond electron density and the bottom image the molecular dipole

m = 1.02 D

Tetrachloromethane

The top image show the bond electron density and the bottom image the molecular dipole

m = 0 D

Dipoles and Intermolecular Attraction

Melting points and boiling points are important physical properties. These properties reveal something about the forces that hold molecules together in condensed phases (liquids and solids). Chemists recognize three major kinds of attractive forces in covalent molecules, all of which are related to dipoles.

Polar molecules have a permanent dipole moment. Since opposite charges attract, when polar molecules approach each other they orient themselves in a head-to-tail manner. The following example shows the dipole-dipole attraction in chloroform (trichloromethane, bp 61oC).

dpdp.gif (4738 bytes)

Carbon tetrachloride (tetrachloromethane, bp 77oC) is a nonpolar molecule but it is a liquid at room temperature, indicating that some attraction between molecules must exist. The molecule has no permanent dipole but an instantaneous dipole is formed when two CCl4 molecules approach each other. The electron cloud in one molecule repulses the electrons in the second molecule breaking the symmetry. These temporary dipoles exist for only a short time and fluctuate from one molecule to another. The result is a weak dipole-dipole attraction called the London dispersive force (van der Waals force). The more contact area between molecules the stronger the van der Waals forces. We will see examples of this trend when we examine the boiling points of hydrocarbons.

 

vdw-1.gif (2867 bytes) vdw-2.gif (4008 bytes)
Tetrachloromethane molecules far apart.  No dipole moment. Induced dipole moment of two tetrachloromethane molecules close together.

Hydrogen bonding is the result of strong dipole-dipole attraction in molecules containing O-H and N-H bonds. HF also undergoes hydrogen bonding but since F is monovalent, this bond is not found in organic molecules. These hydrogen bonds are strongly polarized with the hydrogen atom carrying a partial positive charge. Although hydrogen bonding is only about 10% as strong as covalent bonds it is responsible for the unusual high boiling points of water and alcohols which contain O-H bonds. Ammonia and amines contain N-H bonds which are less polar than O-H bonds and the resulting hydrogen bonding is weaker.