Bonding Theories | General Chemistry 1

The covalent bonding is studied in this chapter: valence bond theory, molecular orbital theory, multiple bonds, hybridization and hybrid orbitals, molecular orbital diagrams.

Valence Bond Theory

The Lewis structures and the VSEPR theory are used to predict molecular geometry but do not tell us anything about bond formation. There are 2 theories of bonding:

  • Valence bond theory ⇒ bonds are formed when atomic orbitals overlap
  • Molecular orbital theory ⇒ quantum mechanical treatment of bonding


Valence bond theory:

Atoms share valence electrons when an atomic orbital (AO) on one atom overlaps with an atomic orbital on the other. A bond forms when the resulting molecule is lower in energy than the original free atoms
 

σ bond vs. π bond

σ bond results from an overlap along the internuclear axis: 2s orbitals, 1s-1p orbitals, or 2p orbitals
π bond results from an overlap located above and below the internuclear axis: side-by-side overlap of 2p orbitals
 

Multiple Bonds and Hybridization

Single bond: all single bonds are σ-bonds (free rotation around σ-bonds)
Double bond: 1 σ-bond + 1 π-bond (restricted motion around π-bonds)
Triple bond: 1 σ-bond + 2 π-bonds

 

H2 ⇒ 1 σ-bond
O2 ⇒ 1 σ-bond + 1 π-bond
N2 ⇒ 1 σ-bond + 2 π-bonds

 

 

It is difficult to explain the shapes of simple molecules only with atomic orbitals:
atomic orbitals combine to give hybrid orbitals ⇒ hybridization.
Hybridization does not occur when atoms are not bonding.

Number of hybrid orbitals formed = number of atomic orbitals combined

 

1 s-orbital + 1 p-orbital can combine ⇒ 2 equivalent sp hybrid orbitals are formed
1 s-orbital + 2 p-orbitals can combine ⇒ 3 equivalent sp2 orbitals are formed
1 s-orbital + 3 p-orbitals can combine ⇒ 3 equivalent sp3 orbitals are formed

Hybrid Orbitals

sp hybridization: 1 s-orbital + 1 p-orbital
⇒ 2 equivalent sp atomic orbitals ⇒ 2 electron group ⇒ linear

sp2 hybridization: 1 s-orbital + 2 p-orbitals
⇒ 3 equivalent sp2 atomic orbitals ⇒ 3 electron group ⇒ trigonal planar

sp3 hybridization: 1 s-orbital + 3 p-orbitals
⇒ 4 equivalent sp3 atomic orbitals ​​​​​​​⇒ 4 electron group ​​​​​​​⇒ tetrahedral

 

 

 

Hybrid orbitals from d orbitals:

sp3d hybridization: 1 s-orbital + 3 p-orbitals + 1 d-orbital
​​​​​​​⇒ 5 equivalent sp3d atomic orbitals ⇒ 5 electron group ⇒ trigonal bipyramidal

sp3d2 hybridization: 1 s-orbital + 3 p-orbitals + 2 d-orbital
⇒ 6 equivalent sp3d2 atomic orbitals ⇒ 6 electron group ⇒ octahedral

 

sp3d hybridization: orbitals of PCl5
​​​​​​​
sp3d2 hybridization: orbitals of SF6

Molecular Orbital Theory

This theory explains: bond formations + magnetic and spectral properties of molecules.
It assumes that molecules have orbitals (molecular orbitals MO) in the same way that atoms have orbitals (atomic orbitals AO).


Atomic wave functions are summed to obtain molecular wave functions:

  • If wave functions reinforce each other ⇒ bonding MO formation
  • If wave functions cancel each other ⇒ antibonding MO formation (with a node)

Bonding MO is lower in energy than the AOs ⇒ more stable
Antibonding MO is higher in energy than the AOs ⇒ less stable


2 overlapping s-orbitals form 1 σ bonding MO + 1 σ* antibonding MO
3 overlapping p-orbitals in a p subshell form:

  • 1 σ bonding MO + 1 σ* antibonding MO
  • 2 π bonding MOs + 2 π* antibonding MOs

 

Molecular Orbital Diagrams

Molecular orbital diagrams show the relative energy and number of electrons in each MO + the AOs they are formed from. The MO with the lowest energy is filled first.

 

 

 

Bond order = 1/2 [(number of e- in bonding orbitals) – (number of e- in antibonding orbitals)]
The higher the bond order, the stronger the bond

 

MO energy levels for O2, F2 and Ne2:

 

MO energy levels for B2, C2 and N2: