What is the Lewis structure for the molecule O3? What is formal charge and resonance? Based on electronegativity differences, would you expect the O-O bonds in ozone to be polar? It turns out that these bonds are polar; what could explain this?

The Lewis structure of ozone (O₃) can be represented as a resonance hybrid of two structures. In each structure, one double bond and one single bond are present between the oxygen atoms. The Lewis structure can be depicted as follows:

   O
  // \
 O    O
  \ // 
   O

In this diagram, the central oxygen atom is bonded to two other oxygen atoms: one through a double bond and one through a single bond. Each oxygen atom, except for one, bears a formal charge. To calculate the formal charge, we use the formula:

• Formal Charge = (Valence Electrons) – (Non-bonding Electrons) – 0.5*(Bonding Electrons)

Each oxygen atom typically has six valence electrons. In the ozone molecule’s most stable resonance structures, the ozone molecule has a formal charge distribution of +1 on one oxygen and -1 on the other, with the central oxygen atom having a formal charge of 0. This distribution leads to the conclusion that ozone is a resonance hybrid rather than a fixed structure.

Now, regarding the polarity of the O-O bonds in ozone, it is important to consider the electronegativity of the oxygen atoms. Oxygen has an electronegativity value of 3.5 on the Pauling scale. However, since both oxygen atoms have the same electronegativity, one may initially think that the O-O bonds would be non-polar. Nonetheless, the presence of a formal charge and the resonance structure introduces a slight difference in electron density between the bonds, leading to a dipole moment that makes the O-O bonds in ozone polar.

This polarization occurs because the bond lengths and character are not equivalent in the resonance forms, meaning that despite identical electronegativity, the interaction of the surrounding electron cloud leads to distribution effects that create slight dipoles.