VALENCE BOND THEORY & VSEPR:
Valence bond theory and Valence Shell Electron Pair Repulsion theory (VSEPR) are simple
but effective models for understanding the bonding that occurs in organic compounds.
 Covalent bonds are usually formed between two atoms by the overlap of 2 half-filled
(singly-occupied) orbitals of the bonding atoms. For example, two hydrogen atoms form
a covalent bond by the overlap of their half-filled 1s atomic orbitals.
H
1s1
+
H
H
H
H
H
1s1
Valence Bond theory predicts that
two, half-filled 1s atomic orbitals
of hydrogen overlap to form a single
(sigma) bond.
This H 2 molecule is stable (low
in energy) because the electrons
spend most of their time between
H nuclei-drawing them together.
Each H atom in
H2 is isoelectronic
with Helium
 Carbon is in group IVA of the periodic table and therefore has 4 valence (bonding)
electrons, i.e., 4 electrons in its highest Principal Energy Level (PEL). Its ground state
electron configuration is given as …
1s2 2s2 2px1 2py1 2pz0
4 valence electrons
 Carbon normally forms 4 covalent bonds using its 4 valence shell electrons, however, its
ground state electron configuration is not suitable for this because the 2s orbital is full and
the 2pz orbital is empty. In a process called ‘hybridization’, one of the 2s electrons is
promoted to the vacant 2pz orbital producing 4 half-filled orbitals. These orbitals alter
their shape in order to achieve the best possible overlap with orbitals from other atoms
and also to move as far away from each other as possible.
1s2 2s2 2px1 2py1 2pz0
hybridization
1s2 2s1 2px1 2py1 2pz1
 Only the outer, valence electrons are involved in bonding and thus the inner shells of
electrons will not be considered further.
 The vast majority of organic compounds are formed from just a few elements, i.e., C, N,
O, P, S and the halogens. We will look at their hybridization states in the following pages.
ORGANIC CHEMISTRY INTRO
1
BONDING IN CARBON COMPOUNDS
6
C
orbital shape
Group 4A
4 valence e’s.
Covalence = 4
2py1
2px1
2s2
ground state
valence electron
configuration
A change (hybridization) occurs
to the orbital shape and electronic
configuration to facilitate bonding.
The 3 hybridizations occurring in
cabon are shown below.
2pz0
sp 2 hybridized
(triangular planar)
sp hybridized
(linear)
Carbon forms a double bond
Carbon forms 1 triple or 2 double bonds
sp 3 hybridized (tetrahedral)
Carbon forms 4 single bonds
The arrangement of the 4 atomic
orbitals and the electronic
configuration in the carbon atom
are not suitable to form 4 bonds.
180º
90º
109.5º
or
C
120º
C
C
C
orbital
shape
(sp2 orbitals x 3) + (p x 1)
hybridized
orbitals
(sp orbitals x 2) + (p x 2)
(sp3 orbitals x 4)
electron
configuration
(3 s bonds) + (1 p bond)
(forms 4 s bonds)
alkanes
e.g., ethane (C 2H6)
bonds
C
C
H
H
H
H
H
C
C
H
H
H
C
H
C
H
H
H
H
H
ORGANIC CHEMISTRY INTRO
H
H
C
H
H
H
H
C
C
H
H
H
H
H
H
H
H
H
C
alkynes
e.g., acetylene (C 2H2)
H
H
H
H
(2 s bonds) + (2 p bonds)
alkenes and arenes
e.g., ethene and benzene
(C 2H4)
(C 6H6)
H
H
C
C
H
also CO2
O C O
2
BONDING IN NITROGEN COMPOUNDS
Group 5A
7
The shape and orientation of the
4 atomic orbitals in the nitrogen
atom are not suitable for forming
3 (or 4) bonds.
N
orbital shape
5 valence e’s.
Covalence = 3
2s2
ground state
valence electron
configuration
2py1
2px1
A change (hybridization) occurs
to the orbital shape to facilitate
bonding. The 3 hybridizations
occurring in nitrogen are shown
below.
2pz1
sp 3 hybridized
sp 2 hybridized
sp hybridized
Nitrogen forms 3 or 4 single bonds
Nitrogen forms a double bond
Nitrogen forms 1 triple bond
......
..
..
90º
or
.:.
N
N
107º
orbital
shape
hybridized
orbitals
180º
N
120º
..
....
..
(sp 2 orbitals x 3) + (p x 1)
(sp3 orbitals x 4)
(sp orbitals x 2) + (p x 2)
electron
configuration
bonds
N
..
1 lone pair
+
3 s bonds
....
or
H
N
N
H
H
(1 lone pair) + (2 s bonds) + (1 p bond)
(4 s bonds)
Cl H
ammonium
chloride
NH3
NH4 Cl
+ -
nitriles
e.g., ethanenitrile
N
CH3
C
..:
N
(1 lone pair) + (1 s bond) + (2 p bonds)
..
N
amines
ammonia
ORGANIC CHEMISTRY INTRO
..
+
H
H
H
azo compounds
e.g., (trans azobenzene)
.:.
N
....
N
:N
C
CH3
3
BONDING IN OXYGEN COMPOUNDS
8
The shape and orientation of the
4 atomic orbitals in the oxygen
atom are not optimal for forming
2 (or 3) bonds.
O
orbital shape
Group 6A
6 valence e’s.
Covalence = 2
2s2
ground state
valence electron
configuration
2px2
2py1
A change (hybridization) occurs
to the orbital shapes to facilitate
bonding. The 2 hybridizations
occurring in oxygen are shown
below.
2pz1
sp 3 hybridized
sp 2 hybridized
Oxygen forms 2 single bonds
Oxygen forms a double bond
O
..
..
hybridized
orbitals
.:.
.... O
....
..
orbital
shape
90º
105º
or
..
..
O
:
120º
..
(sp 2 orbitals x 3) + (p x 1)
(sp3 orbitals x 4)
electron
configuration
bonds
2 lone pair
+
2 s bonds
or
e.g., methanol (CH 3OH)
CH3
.... O
....
e.g., acetone (CH3CCH3)
+
hydronium ion (H3O )
H
.. O
H
ORGANIC CHEMISTRY INTRO
(2 lone pair) + (1 s bonds) + (1 p bond)
O
1 lone pair
+
3 s bonds
H
..:
O
..:
C
+
H
H3C
: O:
C
H3C
CH3
CH3
4
 Note that a double bond is made of one s and one p bond.
 Note that a triple bond is made of one s and two p bonds.
 Halogens (groups VIIA elements) generally form only one s bond in organic compounds.
They do not reshape their orbitals (hybridize) when they bond.
The shape and orientation of the
4 atomic orbitals in the halogens
are adequate for forming one
single bond.
17
Cl
orbital shape
ground state
3s2
valence electron
configuration
3px2
3py2
(3 lone pairs and 1 s bond)
H
Cl
Hybridization does not occur
when halogens form single
bonds.
3pz1
C
H
H
..
: Cl
..
CH3
methyl chloride
(chloromethane)

Hydrogen, like the halogens, does not hybridize its 1s orbital when bonding.

Silicon, like carbon, is a group 4A element with 4 valence electrons. As expected, silicon forms
sp3 hybridized tetrahedral compounds with 4 substituents. Simple examples include silicon
tetrabromide (SiBr4) and tetramethylsilane [(CH3)4Si]. Silicon forms a few compounds in which
it has double bonds, e.g., H2Si=CH2. However, silicon's large size makes p-orbital overlap for p
bonds less effective than in carbon compounds. Unlike 2nd period elements which cannot
accommodate more than 8 electrons in their valence orbitals, Si, a 3rd period element can expand
its valence shell to accommodate 10 electrons (sp3d hybridized – 5 substituents, e.g., SiF5-) and
even 12 electrons (sp3d2 hybridized – 6 substituents, e.g., fluorosilicic acid, H2SiF6).

Phosphorus, like nitrogen, is a group 5A element with 5 valence electrons. As expected,
phosphorus forms sp3 hybridized compounds with 3 substituents. Simple examples include
phosphorus tribromide (PBr3) and trimethylphosphine [(CH3)3P]. Phosphorus forms some
compounds in which it has double bonds to oxygen, e.g., phosphoric acid (H3PO4). However,
phosphorus’ large size makes p-orbital overlap for p bonds less effective than in C or N
compounds. Like other 3rd period elements, phosphorus can be bonded to 4, 5, and 6 atoms.
e.g., phosphorus oxychloride (Cl3P=O), phosphorus dibromide trichloride (PBr2Cl3), and
phosphorus hexafluoride anion (PF6-).

Sulfur, like oxygen, is a group 6A element with 6 valence electrons. As expected, sulfur forms
sp3 hybridized compounds with 2 substituents. Simple examples include dimethyl sulfide
[(CH3)2S] and methyl mercaptan (methane thiol) (CH3SH). Sulfur can form bonds to three
(H2SO3), four (H2SO4), five (SOF4) and six atoms (SF6).
 A methyl cation has an sp2 hybridized carbocation with a vacant p orbital. A methyl radical has
an sp2 hybridized carbon atom with a ½-filled 2p orbital. A methyl anion contains an sp3
hybridized carbanion with a lone pair in one of its sp3 orbitals. Draw them.
ORGANIC CHEMISTRY INTRO
27
In saturated compounds, all atoms have only s bonds, whereas in unsaturated compounds,
one or more p bonds are present.
 Conjugated unsaturation occurs when alternating s and p bonds are present. In such
compounds, all p-orbitals in conjugated p bonds overlap.
1,3-butadiene
CH2
CH
H
CH2
H
H
O
C
C
C
CH
H
C
H
H
H
H
H
C
C
C
2-cylcopentenone
H
C
H
H
 Isolated unsaturation occurs when p bonds are separated by more than one s bond.
In such compounds, p-orbitals of one p bond cannot overlap with p-orbitals of other p
bonds.
CH2
1,4-pentadiene
CH
CH2
H
O
C
H
C
C
H
C
H
H
H
H
H
H
3-cylcopentenone
C
C
H
C
C
H
CH2
H
H
C
CH
H
C
H
H
 Cumulated unsaturation describes immediately adjacent unsaturation. Cumulated
carbon-to-carbon compounds are not very stable and are rarely encountered.
1,2-butadiene
ORGANIC CHEMISTRY INTRO
H2C
C
CH
CH3
28
HYBRIDIZATION STATE IS BASED ON THE NUMBER OF REGIONS OF ELECTRON DENSITY
Be
sp
H
Be
C
B
..
O
..
H
N
..
O
..
C
C
H
C
H
H
C
H
H
C
H
..
HO
..
..
O
..
..
O
..
..
O H
..
-
H
+
N
C
B
H
H
..
OH
..
..
O
H
H
H
C
H
H
H
..
O
N
H
H
H
H
H
NH4+Cl-
Si
+
+
H
Al
:
H
H
H
Na+BH4-
Mg
:
..
H
O
H
H
sp3
:
-: .. :
O
N
-
H
..
F
:
..
N
..
N
+
sodium azide
H
..
O.. H
..
O
..
Na+
N
N
H
H
C
..
N
..
C
B
sp2
H
O
hydronium ion
P
Cl
S
sp
H
H
H
Al
sp2
..
H
P
C Si
H
H
H
CH2
:O :
:O :
..
S
H3C
S
:O :
:O:
:O :
:O :
:O :
S
:O :
-
H
OH
sp3
HO
HO
Al
H
H
H
+
Li
AlH4-
HO
HO
Si
HO
HO
P
Cl
..
..O
HO
:O :
..
Cl
S
P
Cl
..
O
..
:O :
OH
OH
Cl
..
Cl :
..
HO
HO
Cl
Cl
ORGANIC CHEMISTRY INTRO
..
.O
.
..
..O
:
:
HO
:O :
..
Cl
S
:
:
HO
29
:
Study the following table. In the last 3 columns Lewis structures are drawn as if the atoms
were bonded. Learn these names and structures and identify their hybridization states.
Lewis
Symbol
#
valence
#
#
bonds
unshared
e- 's
e- 's
+ 1 F.C.
neutral
- 1 F.C.
B
B
3
3
0
-
B
none
boride ion
..
C
+
C-
C
C
4
4
0
carbonium
ion
N
N
5
3
2
carbide ion
2
4
N
N
:
..
..
O
:
F
7
1
6
Cl +
:
7
1
6
..
.. F
..
..
F
..
:
..
fluoronium ion
Cl
..
an oxide ion
..
F+
O-
:
oxonium
ion
:
-
nitride ion
O+
6
..
nitronium
ion
..
O
..
+
..
chloronium ion
unhybridized
..
..
.. Cl
unhybridized
unhybridized
fluoride ion
.. ..
Cl :
..
unhybridized
chloride ion
Bromine and iodine are analogous to fluorine and chlorine. Draw the structures of
bromonium and iodinium cations, bromide and iodide anions, and bromine and iodine.
ORGANIC CHEMISTRY INTRO
30