CH221 CLASS 3
CHAPTER 2: POLAR COVALENT BONDS; ACIDS AND BASES
Synopsis. The third class deals with ideas about polar covalent bonds,
electronegativity, dipole moments, formal atomic charges and resonance.
My name is Bond, Covalent Bond: shared not taken
Polar Covalent Bonds and Electronegativity
The two major types of chemical bonding: ionic bonding, involving complete
electron transfer, and covalent bonding, involving sharing of electrons, are in
reality descriptions of two extreme circumstances. The electronic nature of most
chemical bonds lies somewhere between these two extremes. In particular, many
bonds in organic molecules may be described as polar covalent bonds, indicating
that the bond electrons are not shared equally (symmetrically), meaning that the
electron distribution in the bond is not symmetrical.
Bond polarity occurs when there is a difference in electronegativity between the
bonded atoms. Electronegativity is the term used to describe the ability of a
bonded atom to attract bond electrons: it can be can be measured on a scale
(devised by L Pauling and R Mulliken) that arbitrarily sets the electronegativity of
hydrogen as 2.1. All the other elements then take up values between 0.7 (Cs)
and 4.0 (F), as shown on the next page.
Group
Period 1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
1
H
2.1
2
Li Be
0.98 1.57
B
C
N
O
F
Ne
2.04 2.55 3.04 3.44 3.98 0
Na
Al
3
He
0
Mg
0.93 1.31
Si
P
S
Cl
Ar
1.61 1.9 2.19 2.58 3.16 0
4
K
Ca
0.82 1
Sc Ti V
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
1.36 1.54 1.63 1.66 1.55 1.83 1.88 1.91 1.9 1.65 1.81 2.01 2.18 2.55 2.96 0
5
Rb Sr Y
Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I
Xe
0.82 0.95 1.22 1.33 1.6 2.16 1.9 2.2 2.28 2.2 1.93 1.69 1.78 1.96 2.05 2.1 2.66 2.6
6
Cs Ba La Hf Ta W Re Os Ir
Pt Au Hg Tl Pb Bi Po
0.79 0.89 1.1 1.3 1.5 2.36 1.9 2.2 2.2 2.28 2.54 2
2.04 2.33 2.02 2
7
Fr Ra Ac
Rf
0.7 0.89 1.1
Lanthanides
Actinides
Db
Sg
Bh Hs
Mt
At Rn
2.2 0
Uun Uuu Uub
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
1.12 1.13 1.14 1.13 1.17 1.2 1.2 1.1 1.22 1.23 1.24 1.25 1.1 1.27
Th Pa U
Np Pu Am Cm Bk Cf Es Fm Md No
Lr
1.3 1.5 1.38 1.36 1.28 1.3 1.3 1.3 1.3 1.3 1.3 1.3 1.3
Key: Electronegativity
No 0-.66 .66-1 1-1.33 1.33-1.66 1.66-2 2-2.33 2.33-2.66 2.66data
Polar covalent bonds have tiny dipole moments arising from the charge
separation in the bond, as illustrated below

H3C

F
 
H3C Li
Electronegativity of
carbon is less than
that of fluorine, but
greater than that of
lithium
The ability of an atom to polarize a bond is often called the inductive (I) effect.
Atoms of low electronegativity, like Li (above) and Mg are electron donors (+I),
whereas atoms of high electronegativity, like N, O and F, are electron
withdrawers (-I).
The inductive effects in polar covalent bonds play a major part in
the understanding of chemical reactivity and will be much used
in future.
Polar Covalent Molecules: Dipole Moments
The dipole moment () of a molecule is a useful measure of the polar character
of the molecule and can be measured in the laboratory, using electromagnetic or
spectroscopic methods. It can be regarded as the vector sum of the individual
bond dipole moments, which are discussed next.
Definition of dipole moment: if two charges of equal magnitude and opposite sign
(i.e. +Q and –Q) are separated by a distance R, then
Dipole moment ()
It is symbolized as shown above
=
QR
Units of Dipole Moment and Estimation of Bond Polarity
The SI unit of  is coulomb.meters, but for atomic systems, a more useful unit is
the debye (D), which is equal to 3.336 x 10-30 coulomb.meters.
1D is defined as the dipole moment of two charges +e and –e that are 0.2082Å
apart. (1Å = 100 pm). Therefore, if  is the fraction of charge ( e) on each atom
of a bond, which are R (Å) apart (i.e. Q = e), then
 = R
0.2082
E.g. For CH3Cl, the measured dipole moment is 1.87D and RC-Cl is
known to be 178 pm (= 1.78 Å). Assume the total contribution to
the measured dipole moment comes from the C-Cl bond.
Then,  =
or
=
1.87
=
1.78
0.2082
0.22
That means that the fraction of unit charge on carbon is +0.22e
and on chlorine is –0.22e: the C-Cl bond is 22% ionic.
The molecular dipole moment, being the vector sum of the bond dipole moments,
is very much dependent upon molecular geometry, as illustrated by the examples
below.
Net
.. H
:O
H
 = 1.85D
H
Net
:N
H
H
 = 1.47D
Cl
C
Cl
Cl
Cl
Net: zero
O=C=O
Net: zero
 = 0D
 = 0D
The two lower molecules have symmetrical structures and hence the individual
bond polarities exactly cancel.
Class Question
Make three-dimensional drawings of the following molecules and predict which
have dipole moments. Show the direction of the dipole moment.
(a) 1,1-Dichloroethene (b) trans-1,2-Dichloroethene (c) cis-1,2-Dichloroethene (d)
Tetranitromethane (e) Trichloromethane (chloroform) (f) Carbon oxysulfide (OCS)
(g) Sulfur dioxide (h) Ethylene
Formal Charges in Polar Molecules
The assignment of formal charges to specific atoms in a molecule is closely
related to the previously discussed ideas of bond polarity and dipole moment.
Formal charges are most important for molecules that contain atoms with an
abnormal number of bonds (e.g. 3 bonds, instead of 4 bonds for carbon, 4 bonds
instead of 3 bonds for nitrogen, etc). Formal charge is defined below.
Some examples are given on the next page.
Class Question
Determine the formal charges for the non-hydrogen atoms in the following
molecules.
(a) CH2=N=N (diazomethane) (b) (CH3)3N-O(trimethylamine oxide)
(c) CH3-NC (methyl isocyanide (d) Ph-N=N=N (phenyl azide)
Examples
1. No "abnormal" valence
H
Chloromethane
..
H :C: H
..
FC = 4 - 4 - 0 = 0
:Cl:
..
H:N:H
..
H :C: H
..
H
FC = 7 - 1 - 6 = 0
FC = 5 - 3 - 2 = 0
FC = 4 - 4 - 0 = 0
Methylamine
2. "Abnormal" valence present
O
Nitromethane
FC = 6 - 2 - 4 = 0
H
O
..
H:C:N
..
H O
Dimethyl sulfoxide
FC = 5 - 4 - 0 = +1
S
N
FC = 6 - 6 - 1 = -1
)
O
FC = 6 - 6 - 1 = -1
O
CH3
(i.e. CH3
(i.e. (CH3)2S
O)
FC = 6 - 3 - 2 = +1
CH3
Note that expanded valence (via use of d orbitals) is allowed for period 3
elements and beyond, so that dimethyl sulfoxide can also be written (CH3)2S=O
A summary of formal charges for C, N and O is given in Table 2.2, on p. 36 of the
textbook.
Resonance
For many molecules, a single Lewis structure or line-bond structure does not fully
represent the actual structures of the molecules, as determined by experiment,
notably X-ray, electron and neutron diffraction methods. These molecules, which
are illustrated by conjugated unsaturated groups or aromatic rings, can only be
satisfactorily represented by a number of line-bond structures, each of which
contributes a certain amount to the overall (real) structure. This is demonstrated
for nitromethane, overleaf.
O
O
N
CH3
CH3
N
O
O
O
Overall
CH3
N
O
Measured rNO = 122 pm, for both bonds
rN=O = 116 pm
rN-O = 130 pm
The two conventional line-bond structures above differ only in electron
distribution (position of double bond and lone pairs): they are known as
resonance structures and the phenomenon is called resonance. The overall
(actual) structure can be considered to be formed by the combination (in this
case equal combination) of the two resonance structures and is known as a
resonance hybrid.
The actual structure, represented by the resonance hybrid, is more stable (of
lower energy) than any of the individual resonance structures, and hence
resonance leads to stabilization. Other important examples include benzene and
carboxylate ions.
Other,
less favorable
resonance
structures are
not shown
Measured rCC = 138 pm, for all
CC
rC=C = 134 pm
bonds
rC-C = 154 pm
Overall
O
O
R
C
R
O
Overall
O
R
C
O
Measured rCO= 127 pm
C
O
rC=O = 120 pm
rC-O = 135 pm
for both
The Rules of Resonance
1. Whenever a molecule can be represented by two or more structures that
differ ONLY in the arrangement of electrons, there is said to be
resonance. The individual structures are called RESONANCE
STRUCTURES. The actual structure of the molecule is a RESONANCE
HYBRID of those structures and cannot be satisfactorily represented by
any one of them. Instead, each resonance structure contributes a certain
amount to the overall structure of the molecule.
2. The actual structure of the molecule (represented by the resonance
hybrid) is more stable than any of the contributing resonance structures.
The energy difference is known as the RESONANCE ENERGY and the
more nearly equal in stability of resonance structures, the greater the
resonance energy.
3. Resonance is important when resonance structures are of about equal
stability (i.e. of about the same energy) as in the case of nitromethane.
The amount of contribution of each resonance structure depends on the
relative stability of that structure. Generally, the greater the number of
charges, the less stable is the structure. This point is illustrated in 2,
above. Also structures with like charges on adjacent atoms are unstable,
as are those that require redistribution of electrons in the reverse
direction of electronegativity difference. An example of the latter is
>C=O
major

>C+-Ominor

>C--O+ (insignificant)
4. Stabilization can be achieved in saturated molecules by
HYPERCONJUGATION or “NO-BOND RESONANCE”.
E.g. the ethyl radical
H
H
C
H
H
.
CH2
H
C
CH2
H
.
C
H
major
all minor
OR
H
C
C
CH2
H
C
.
H
H
H
H
H
.
H
H
H
CH2