In this example, you'll learn how to determine whether a molecule has dipole-dipole forces. The problem reads, which molecules have dipole-dipole forces? A.
Carbon dioxide, B. Dichloromethane, and C. Methane. A molecule has dipole-dipole forces if it is polar. To determine if a molecule is polar, 1. determine if the molecule contains polar bonds, and 2. consider the geometry of the molecule to determine if the polar bonds add together to form a net dipole moment.
Part A, carbon dioxide. Since the electronegativity of carbon is 2.5 and that of oxygen is 3.5, carbon dioxide has polar bonds. The geometry of carbon dioxide is linear.
You can determine this from its Lewis structure. As a result, the two polar bonds point in exactly opposing directions and they cancel each other out so that the molecule is not polar and therefore does not have dipole-dipole forces. Part B, dichloromethane. The electronegativity of carbon is 2.5. that of hydrogen is 2.1 and that of chlorine is 3.0.
Consequently, the dichloromethane has two polar bonds, the two carbon-chlorine bonds, and two bonds that are nearly nonpolar, the two carbon-hydrogen bonds. The geometry of CH2Cl2 is tetrahedral. Because the carbon-chlorine bonds and the carbon-hydrogen bonds have different polarities, their dipole moments do not cancel each other, but sum to a net dipole moment.
The molecule is polar and has dipole-dipole forces. C. Methane.
The electronegativity of carbon is 2.5 and that of hydrogen is 2.1. So the carbon-hydrogen bonds are nearly nonpolar. The geometry of the molecule is tetrahedral. Go ahead and answer this one yourself. Does methane have dipole forces?
A. Yes. B. No. The correct answer is B.
No. Since the geometry of the molecule is tetrahedral, any slight polarities that the bonds have cancel. Methane is therefore nonpolar and does not have dipole-dipole forces.