4.7 Sigma and pi bonds (HL only)
Worked example 4.3
2– O1 O4 S O2 O3 Structure 1
Two alternative Lewis structures for SO42− are shown here. Use the concept of formal charge to deduce which is the better representation of the actual structure. Structure 1
2– O1 O4 S O2 O3 Structure 2
Structure 2 Formal Charges
S O1 O2 O3 O4
1 × 8 0 = 2+ 2 1 FC = 6 − × 2 6 = 1− 2 1 FC = 6 − × 2 6 = 1− 2 1 FC = 6 − × 2 − 6 = 1− 2 1 FC = 6 − × 2 − 6 = 1− 2
FC = 6 −
1 × 12 − 0 = 0 2 1 FC = 6 − × 2 6 = 1− 2 1 FC = 6 − × 4 4 = 0 2 1 FC = 6 − × 2 − 6 = 1− 2 1 FC = 6 − × 4 4 = 0 2
FC = 6 −
S O1 O2 O3 O4
The sum of the formal charges on all the atoms should add up to the overall charge on the molecule/ion.
Structure 2 has lower individual formal charges (S in Structure 1 has a 2+ charge) and therefore it is the preferred Lewis structure. T E S T Y O UR SELF 4. 7
Use the concept of formal charge to work out the preferred Lewis structures for: − − SO3 PO43 ClO4 XeO4
4.7 Sigma and pi bonds (HL only) Sigma (σ) bonds result from the axial (head-on) overlap of atomic orbitals. The electron density in a σ bond lies mostly along the axis joining the two nuclei.
p orbital
p orbital
σ
A pi (π) bond is formed by the sideways overlap of parallel p orbitals. The electron density in the π bond lies above and below the internuclear axis (Figure 4.18). Figure 4.17
Single bond
σ
H H
Double bond Triple bond
σ
C σ
H
H
C
σ σ
H
H
σ+π
σ + 2π
σ
σ
H H
σ
σ
C σ
C
π σ π σ π
σ
H
C
C
σ
H H
H
σ σ
σ
C
σ
C
side-on overlap
H
σ
head-on overlap
side-on overlap
H
H σ
sigma bond
σ
H H
H
π C
C
H
H H
π- bond
Figure 4.18 σ
H
41