The sulfur dioxide (SO2) molecule has several resonance structures that help explain its bonding characteristics. To understand these structures, we need to consider the arrangement of electrons and the formal charges on each atom.
Firstly, the Lewis structure of SO2 shows that sulfur is the central atom with two oxygen atoms bonded to it. The basic structure can be drawn as follows:
O
||
S = O
In this structure, sulfur shares double bonds with both oxygen atoms. However, we can also draw additional resonance structures to represent differing electron distributions:
Resonance Structure 1:
O || S | O
In this structure, sulfur has a double bond with one oxygen and a single bond with the other oxygen, with a lone pair of electrons on the singly bonded oxygen atom. The formal charges for this structure are:
- Sulfur (S): 6 – (0 + 1) = 5 (formal charge = +1)
- Oxygen (O with double bond): 6 – (4 + 0) = 2 (formal charge = 0)
- Oxygen (O with single bond): 6 – (6 + 1) = -1 (formal charge = -1)
Resonance Structure 2:
O
|
S
||
O
This structure is the reverse of the first one, with the single bond now on the previously doubly bonded oxygen and vice versa. The formal charges are the same as in the first structure:
- Sulfur (S): 6 – (0 + 1) = 5 (formal charge = +1)
- Oxygen (O with double bond): 6 – (4 + 0) = 2 (formal charge = 0)
- Oxygen (O with single bond): 6 – (6 + 1) = -1 (formal charge = -1)
Both resonance structures contribute to the overall electronic structure of SO2 and the molecule is best represented as a hybrid of these forms. The average formal charges indicate that it is more stable to have no formal charge on the sulfur atom and to have the double bond distributed between the two oxygen atoms.
In conclusion, the resonance structures of SO2 show how sulfur can form bonds with oxygen while still maintaining an overall charge balance in the molecule. Recognizing these structures allows us to better understand the chemical behavior of SO2.