Valence Bond Theory vs. MO Theory
VB Theory begins with two steps:
1) hybridization AOs on atoms participating in bonding
2) Combination of hybrid orbitals to make bonds.
Key differences between MO and VB theory:
1) MO theory: has electrons distributed over whole molecule. (molecule centered)
VB theory: localizes an electron pair between two atoms. (bond centered)
2) MO theory: combines AOs between different atoms to make MOs (LCAO)
VB theory: combines AOs on the same atom to make hybridized atomic
orbitals (hybridization)
3) MO theory: the symmetry must be retained in each orbital.
VB theory: all orbitals must be viewed simultaneously to see retention of the
molecule’s symmetry.
1
2p

2 p unhybridzed
2( sp )  
VB Theory of BeH2
Be
2H
2s 
2px 2py
2px 2py
2 Be-H "antibonds"



sp hybrids

1s 1s
 
Two hybrid atomic orbitals are made to fit the
shape of the molecule, in this case linear, using
atomic orbitals of the Be atom!
2 Be-H bonds
Unused AO are left behind as
unhybridized atomic orbitals
The energy of the hybrid atomic orbitals are intermediate between those of the
original constituent AO’s
The hybrid orbitals combine with other orbitals, atomic or hybrid, creating both
bonding and anti-bonding molecular orbitals, which are localized molecular
orbitals
2p
 
2 p unhybridzed
2( sp 2 )   
VB Theory of BH3
B
2s 
3H
2py
2py
3 B-H "antibonds"




sp2 hybrids


1s 1s 1s
  
3 B-H bonds
BH3 is trigonal planar with three
equal B—H bonds
To get this shape the 2s with two 2p AO’s to generate three equivalent hybrid atomic
orbitals
Combination with the H 1s leads to bonding and anti-bonding molecular orbitals, which
are localized molecular orbitals pointing to the corners of a triangle
2p
  
VB Theory of CH4
no unhybridzed orbitals in 2nd shell
2( sp )    
3
C
4H
2s 
4 C-H "antibonds"





sp3 hybrids



1s 1s 1s 1s
   
4 C-H bonds
CH4 is tetrahedral with 4 equal C-H bonds
To get this shape, we need to combine all the
n=2 AO’s to generate four equivalent hybrid
atomic orbitals
In combination with the H 1s leads to bonding and anti-bonding molecular orbitals
localized and pointing to the corners of a tetrahedron
Valence Bond Theory Summary
Atoms orbitals are hybridized only if it’s necessary to attain the observed
geometry and bond lengths.
Terminal atoms are not typically hybridized.
Geometry determines hybridization:
i) Linear (180º) = sp
(s+p with two leftover p orbitals)
ii) Trigonal planar (120º) = sp2 (s+p+p with one leftover p orbital)
iii) Tetrahedral (109.5º) = sp3
(s+p+p+p with no leftover p orbitals)
Hybrid orbitals combine with each other to make  bonds in which two
electrons are localized between two atoms.
 bonds are made by combining the unused -symmetric p orbitals.
5
Ethane
VSEPR theory requires both carbon atoms to be tetrahedral
The shape of the molecule, requires that contacts be minimized
between the atoms – this is known as the staggered conformation
Bonding is explained by using sp3 hybrid orbitals on each C.
The H atoms bond using their 1s atomic orbitals
In all there are 14 electrons or 7 electron pair bonds in the molecule
1  C-C sp3-sp3 single bond and 6  C-H sp3-s single bonds are formed
H H
H
C
H H




 
C

H
= H 1s
= C sp3
Double bonds: ethene
If we treat ethane by the VSEPR theory, we find that
both carbon atoms are trigonal planar
The molecule is planar. Why ?
sp2 hybrid orbitals on each carbon atom, which leaves
one atomic p orbital unused on each C atom, while H
atoms use their 1s atomic orbitals
There are 6 e’ pair bonds in the molecule, 5 in σ orbitals, 1 in the π orbital
1  sp2-sp2 C-C bond, 1  px-px C-C bond , and 4  sp2-s C-H bonds
The sigma skeleton of ethene
The pi manifold of ethene





= H 1s

= C sp2
Change perspective
to show the π bond!
2 px
2 px
2p
  
no unhybridzed orbitals in 2nd shell
2( sp3 )    
VB Theory of H2O
O
2H
2s 
2 LP’s


 
 
sp 3 hybrids


1s 1s
 
2 C-H
H2O is bent and belongs to the tetrahedral
family with 2 BP and 2 LP
The s and p orbitals combine sp3 hybrids
The 6 e’s from O, singly occupy 2 sp3
orbitals and doubly occupy the remaining
2 as LP’s
The 2 sp3 orbitals combine with 2
1s orbital to form 2 C-H bonds
Triple Bonds: Ethyne
The Lewis structure for ethyne (C2H2)
VSEPR theory: each carbon atom is linear hence sp hybridized
i) C–H  bonds: combination of an sp orbital from C and a 1s orbital from H.
ii) C–C  bonds: combination of sp orbitals from each C.
iii) C–C  bonds: combination of 2p orbitals from each C.
9
Formaldehyde
The Lewis structure for formaldehyde (CH2O) .
VSEPR theory: Carbon atom is trigonal planar hence sp2 hybridized:
i) C–H  bonds: combination of an sp2 orbital from C and a 1s orbital from H.
ii) C–O  bond: combination of an sp2 orbital from C and a sp2 orbital from O.
iii) C–O  bond: combination of 2p z orbitals from C and O.
iv) Lone pairs: remaining sp2 hybrid AOs of O.
10
VB Theory of HCN
The Lewis structure for hydrogen cyanide (HCN) is.
VSEPR theory: C and N atoms are linear hence
are sp hybridized.
i) C–H  bonds: combination of an sp orbital from C
and a 1s orbital from H.
ii) C–N  bond: combination of an sp orbital from C
and a 2p orbital from sp ybrid orbital on N.
iii) C–N  bonds: combination of 2px and 2py orbitals
from C and N.
iv) Lone pair: remaining sp hybrid atomic orbital (2s).
11
Bonding in large molecules
sp2
sp3
sp2
sp3
sp3
sp3
 s sp H -O bond
 sp sp O-C bond
2  s sp H-C bonds
 sp sp C-C bond
 s sp H-C bonds
 sp sp C-N bond
-
sp3
3
3-
-
3
3
3-
-
3
3
3-
3
2
 s sp
-
3
N-H bonds
1 sp3 LP
2-
 sp sp C-C bond
 sp sp C-O bond
 s sp H-O bond
-
3-
2
2-
3
3
 sp sp O-C bond
 p p O-C bond
2
-
2 sp2 LP’s