Organic Chemistry
most electronegative
least elect ronegat iv e
Linus Pauling (1901-1994)
Nobel Prize in Chemistry 1954
Nobel Peace Prize 1962
Electronegativity and Types of Chemical Bonds
The electronegativity difference between interacting
atoms determines the type of bond that is formed.
compound
electronegativity
difference
type of bond
F2
4.0 - 4.0 = 0
nonpolar
covalent
HF
4.0 - 2.1 = 1.9
polar
covalent
LiF
4.0 - 1.0 = 3.0
ionic
Lithium Fluoride
Lithium, a metal, has a very low electronegativity. Fluorine, a
nonmetal, has a very high electronegativity. When they react,
lithium gives up its single outermost (valence-level) electron,
while fluorine takes on a single electron to fill the second energy
level (valence-level).
Li
+ F
Li
+ F
elect ron
configurat ion
of neon
elect ron
configurat ion
of helium
In the solid state, lithum fluoride forms a very stable ionic crystal
lattice structure where each Li+ is surrounded by six F- , and each Fis surrounded by six Li+. Strong electrostatic attractive forces among
the ions of opposite charge stabilize the solid state structure.
Covalent Bonding
When two atoms of the same or similar electronegativities react, they
achieve a noble gas electron configuration by sharing electrons in
covalent bonds. A two electron bond is shown by a dash or line.
Examples
H2
H.
.
. C . + H.
.
H :H or
H-H
H
H
H :C :H or H C H
H
H
: :
CH 4
H. +
Note: Carbon, Group 4, needs 4 electrons to reach the electron
configuration of neon. Hydrogen achieves the electron configuration
of helium by forming a single two electron bond.
Multiple Covalent Bonds and the Octet Rule
Covalent bonds between atoms may involve 4 or 6 electrons.
These are called multiple covalent bonds.
Example N2
Each nitrogen atom has 5 outermost or valence-level electrons.
.
:N. .
The covalent bonding in N2 involves 6 electrons:
N
N
or
:N
N:
The Octet Rule
By sharing 6 electrons, each nitrogen achieves the electron
configuration of neon with an octet of electrons in the valence level.
This tendency to reach 8 electrons in the valence level is called the
octet rule.
Examples of the Octet Rule
N
N
nitrogen
: :
: :
H
H:C :H
H
methane
H:O:H
H :F:
water
hydrogen
fluoride
The above structures are called Lewis structures in honor of G.N.
Lewis. In a Lewis structure, all the valence-level (outermost)
electrons are shown as dots. The Lewis structure of an atom is the
chemical symbol with the valence-level electrons shown as dots.
Example CH3F methyl fluoride
The Lewis structures of the atoms are:
: :
.
. C . .F H . H . H .
.
or
H
HC F
H
: :
H
H
H :C : F : or H C F :
H
H
14 valence electrons
: :
: :
:
The remaining 6 valence
electrons are nonbonding
electrons around the
fluorine atom. The octet
rule applies to both the
carbon and fluorine atoms.
H
H :C : F
H
: :
The four covalent bonds
to the central carbon
atom account for 8 of the
valence electrons:
There are 4 + 7 + 3 = 14
valence electrons available for
chemical bonding and as
nonbonding electron pairs.
Structural Theory of Organic Chemistry
Independent work by August Kekulé, Archibald Couper, and
Alexander Butlerov between 1858 and 1861 laid down the basis of
structural theory in organic chemistry.
(1) The number of bonds to different elements in organic
compounds was deduced. These are valencies.
C
O
N
H
tetravalent
divalent
trivalent
monovalent
(2) Multiple bonds are possible as shown for carbon.
C C
C C
C C
single
double
triple
In all cases, the tetravalency of carbon is maintained.
The Tetrahedral Shape of Methane
In 1874 J.H. Van't Hoff and J.A. Le Bel proposed that the four bonds
to tetravalent carbon point to the corners of a tetrahedron. A regular
four-sided geometric figure.
a tetrahedron
In methane, CH4, each
bond angle is a perfect
109o 28'.
This geometry at carbon
accounts for the overall shapes
of organic structures
Bonding in Simple Hydrocarbons
Hybrid Orbitals and Molecular Shape
sp3 hybridization
CH4
The ground state electron configuration of C is
1s2 2s2 2p2
or
1s
2s
2px
valence level
This electron configuration explains
neither the bonding capacity of C
nor the shape of CH4 .
2py
2pz
Hybrid Orbitals
Step One
Electron Promotion 2s
2p
+energy
1s
2s
2px
2py
2pz
1s
2s
ground state of C
Step Two
1s
2s 2px 2py 2pz
This set of valence
orbitals does not
explain the molecular
shape of methane.
2px 2py 2pz
excited state
Orbital Hybridization
1s
sp3 sp3 sp3 sp3
hybrid orbitals
Overview
orbital hybridization
electron promotion
+ energy
1s
2s 2px 2py 2pz
ground state for C
1s
2s 2px 2py 2pz
excited state
2px 2py 2pz
2s
2s
1s
1s
sp3 sp3 sp3 sp3
hybrid orbitals
sp3 sp3 sp3 sp3
energy
2px 2py 2pz
1s
1s
Mixing of Atomic Orbitals to Make Hybrid Orbitals
Mix 1 2s with 3 2p orbitals
get
4 sp3 orbitals
25% s character 75% p character
Sigma Bonds
Because sp3 orbital has the character of a p orbital, the positive lobe of the
sp3 orbital is large and extends out far from the carbon nucleus.
Sigma Bonds
The positive lobe of the sp3 orbital overlaps with the 1s orbital of hydrogen to
form the bonding molecular orbital of a carbon-hydrogen bond. The two
orbitals have a large overlap due to their size and shape. This large overlap
results in a very strong bond.
The bond formed from the overlap of an sp3 orbital and a 1s orbital is a
sigma (σ) bond. Sigma bonds are bonds in which the orbital overlap gives a
bond that has a circular, symmetric cross section when viewed along the
bond axis. ALL PURELY SINGLE BONDS ARE SIGMA BONDS .
The Shape of CH4
1s
H
1s
H
sp3
sp3
sp3
C
1s
H
sp3
H
1s
H
H
o
109.5
H
H
C
H
H
C
H
H
four sigma bonds
a tetrahedral geometry
Bonding in Ethane
The Structure of Ethane
Ethane, C2H6, is the second member of the alkane family. The four
covalent bonds around each carbon are projected towards the corners of
a tetrahedron, as in methane. This shape of ethane may be predicted by
removing a C-H bond from two methanes and joining the carbons
through a C-C bond.
H
H
H
C
H
H
H
H
H
C
C
H
H
H
HH
H
H
H
H
H
This geometry
follows from the sp3
hybridization at
each carbon.
C
C
C
H
sigma
C-C bond
H
H
The C-C sigma bond results
from the in-phase combination
of sp3 orbitals. Each C-H bond is
formed from the in-phase
combination of an sp3 orbital at
carbon with the hydrogen 1s
orbital.
Rotation Around the C-C Bond
All sigma type bonds have circular symmetry along the bond which means
that there is no loss of orbital overlap when one atom is rotated.
Consequently, there is no significant energy barrier (no increase in
energy) with rotation .
For a C-H bond,
H
there is no
change in energy H
C
with rotation
around the
H
sigma bond.
H
C
H
H
For a C-C bond, there
are small energy
changes with rotation
around the bond that
lead to significant
structural properties.
Formal Charge
Formal Charge
It is possible to assign positive and negative charges to atoms in
Lewis structures. These formal charges often give insight into
chemical reactivity. The sum of all the formal charges must
equal the total charge (if any) on the Lewis structure.
The formal charge on an atom is obtained by subtracting the
number of valence electrons that "belong" to the atom in its bonded
state from the number of valence electrons in the neutral free atom.
Assignment of Valence Electrons in the Bonded State
(1) Electrons in covalent bonds are shared equally by the two
atoms held together by the bond.
(2) Nonbonding electrons are assigned completely to the atom
where they are located.
The total number of valence electrons assigned to an atom in
its bonded state is
one half the number of electrons in covalent
bonds plus all the nonbonding electrons.
Examples of Formal Charges
: :
: :
formal
charge
H
1
-
1
0
C
4
-
4
0
atom
: :
NH3
atom
: :
CH4
H
H C H
H
valence electrons in:
free atom bonded
state
HN H
H
valence electrons in:
free atom bonded
state
formal
charge
H
1
-
1
0
N
5
-
(3 + 2)
0
Examples of Formal Charges
atom
: :
: :
NH4
H
HN H
H
: :
::
: :
CO 32-
formal
charge
H
1
-
1
0
N
5
-
4
+1
atom
:O:
:O:C :O:
valence electrons in:
free atom bonded
state
valence electrons in:
free atom bonded
state
formal
charge
C
4
-
4
0
O
6
-
6
0
O
6
- (6 + 1)
-1
Table of Common Formal Charges
Isomers
Isomers: The Importance of Structure
The development of a structural theory, attention to the linkages among
atoms, solved the problem of isomerism, the existence of different
compounds with the same molecular formula. Isomerism could now be
explained by different linkages.
Example: Isomers of C2H6O
ethyl alcohol
dimethyl ether
boiling point
78.5o C
-24.9o C
melting point
-117.3o C
-138.0o C
reaction with
sodium metal
yes
no
Ethyl alcohol and dimethyl ether are different compounds
with different properties..
Evolution of Chemical Information
Quantitative Analyses
Molecular Formulas
C2H 6O
Structural Theory: Valence
C
C
O
H
H
H
H
H
H
ISOMERS
C
C
H H
O
H
H
H
H
H
H
H
H
H C C O H H C O C H
H H
ethyl alcohol
H
H
dimethyl ether
Ethyl alcohol and dimethyl ether are isomers.
The atoms in ethyl alcohol are connected
differently than those in dimethyl ether. They
are constitutional isomers.