Molecular Polarity

When there are no polar bonds in a molecule, there is no permanent charge difference between one part of the molecule and another, and the molecule is nonpolar. For example, the Cl2 molecule has no polar bonds because the electron charge is identical on both atoms. It is therefore a nonpolar molecule. None of the bonds in hydrocarbon molecules, such as hexane, C6H14, are significantly polar, so hydrocarbons are nonpolar molecular substances.

A molecule can possess polar bonds and still be nonpolar. If the polar bonds are evenly (or symmetrically) distributed, the bond dipoles cancel and do not create a molecular dipole. For example, the three bonds in a molecule of BF3 are significantly polar, but they are symmetrically arranged around the central boron atom. No side of the molecule has more negative or positive charge than another side, and so the molecule is nonpolar:

A water molecule is polar because (1) its O-H bonds are significantly polar, and (2) its bent geometry makes the distribution of those polar bonds asymmetrical. The side of the water molecule containing the more electronegative oxygen atom is partially negative, and the side of the molecule containing the less electronegative hydrogen atoms is partially positive.

 

Sample Study Sheet: Predicting Molecular Polarity

Tip-off – You are asked to predict whether a molecule is polar or nonpolar; or you are asked a question that cannot be answered unless you know whether a molecule is polar or nonpolar. (For example, you are asked to predict the type of attraction holding the particles together in a given liquid or solid.) 

General Steps -

Step 1: Draw a reasonable Lewis structure for the substance.

Step 2: Identify each bond as either polar or nonpolar. (If the difference in electronegativity for the atoms in a bond is greater than 0.4, we consider the bond polar. If the difference in electronegativity is less than 0.4, the bond is essentially nonpolar.)

  • If there are no polar bonds, the molecule is nonpolar.

  • If the molecule has polar bonds, move on to Step 3.

Step 3: If there is only one central atom, examine the electron groups around it.

  • If there are no lone pairs on the central atom, and if all the bonds to the central atom are the same, the molecule is nonpolar. (This shortcut is described more fully in the Example that follows.)

  • If the central atom has at least one polar bond and if the groups bonded to the central atom are not all identical, the molecule is probably polar. Move on to Step 4.

Step 4: Draw a geometric sketch of the molecule.

Step 5: Determine the symmetry of the molecule using the following steps.

  • Describe the polar bonds with arrows pointing toward the more electronegative element. Use the length of the arrow to show the relative polarities of the different bonds. (A greater difference in electronegativity suggests a more polar bond, which is described with a longer arrow.)

  • Decide whether the arrangement of arrows is symmetrical or asymmetrical

  • If the arrangement is symmetrical and the arrows are of equal length, the molecule is nonpolar.

  • If the arrows are of different lengths, and if they do not balance each other, the molecule is polar.

  • If the arrangement is asymmetrical, the molecule is polar.

EXAMPLE – Predicting Molecular Polarity: 

Decide whether the molecules represented by the following formulas are polar or nonpolar. (You may need to draw Lewis structures and geometric sketches to do so.)

a. CO2     b. OF2     c. CCl4     d. CH2Cl2     e. HCN

Solution:

a. The Lewis structure for CO2 is

The electronegativities of carbon and oxygen are 2.55 and 3.44. The 0.89 difference in electronegativity indicates that the C-O bonds are polar, but the symmetrical arrangement of these bonds makes the molecule nonpolar.

If we put arrows into the geometric sketch for CO2, we see that they exactly balance each other, in both direction and magnitude. This shows the symmetry of the bonds.

b. The Lewis structure for OF2 is

The electronegativities of oxygen and fluorine, 3.44 and 3.98, respectively, produce a 0.54 difference that leads us to predict that the O-F bonds are polar. The molecular geometry of OF2 is bent. Such an asymmetrical distribution of polar bonds would produce a polar molecule.

c. The molecular geometry of CCl4 is tetrahedral. Even though the C-Cl bonds are polar, their symmetrical arrangement makes the molecule nonpolar.

                

d.  The Lewis structure for CH2Cl2 is

The electronegativities of hydrogen, carbon, and chlorine are 2.20, 2.55, and 3.16. The 0.35 difference in electronegativity for the H-C bonds tells us that they are essentially nonpolar. The 0.61 difference in electronegativity for the C-Cl bonds shows that they are polar. The following geometric sketches show that the polar bonds are asymmetrically arranged, so the molecule is polar. (Notice that the Lewis structure above incorrectly suggests that the bonds are symmetrically arranged. Keep in mind that Lewis structures often give a false impression of the geometry of the molecules they represent.)

              

e.  The Lewis structure and geometric sketch for HCN are the same:

The electronegativities of hydrogen, carbon, and nitrogen are 2.20, 2.55, and 3.04. The 0.35 difference in electronegativity for the H-C bond shows that it is essentially nonpolar. The 0.49 difference in electronegativity for the C-N bond tells us that it is polar. Molecules with one polar bond are always polar.