Chapter 2
Polar Covalent Bond
Covalent bond in which the electron pairs are not
shared equally.
“Pure”
Covalent
Bond
(non-polar)
X
:
X
Ionic
Bond
increasing bond polarity
X
:Y
X+ Y:-
Electronegativity, χ
ability of an atom in a molecule to attract electron
density to itself.
Structure and Bonding
1
H-F
H-Cl
H-Br
H-I
∆χ
4.0 - 2.1 = 1.9
3.0 - 2.1 = 0.9
2.8 - 2.1 = 0.7
2.5 - 2.1 = 0.4
LiF
BeF2
BF3
CF4
NF3
OF2
F2
Decreasing
polarity
Increasing
ionic
character
∆χ
3.0 ionic
2.4
2.0
1.4
1.0
0.6
0 covalent
Polar covalent bonds: electron density is not shared
equally. Electrons spend “more time” with the more
electronegative atom.
Structure and Bonding
2
So what?
•
Polar bond dipoles are additive and may result in
molecular dipole giving a polar molecule.
Affects physical properties (i.e., melting point, boiling
point, viscosity, vapor pressure).
δ−
δ−
O
F
δ− O
O δ−
C
δ+
non-polar
•
F
δ−
H3C δ+
F
δ−
non-polar
C
CH 3
polar
Regions of high or low electron density provide sites for
chemical reactions to occur.
δ− O
CH3
O
δ−
N
+
δ+ C
H3C
B δ+
H
δ+
H 3C
H3C
R
H δ+
δ+
C
CH3
H 2O +
NH2R
CH3
C
NR
Rules for writing Lewis Structures
Consider Cl2CO (phosgene).
1. Calculate the total number of valence electrons in
the structure. (Don’t forget the charge!)
..
.
:Cl
..
.
.C .
.
Cl
C
O
total
Structure and Bonding
..
.O
.. .
2(7 e -) = 14 e 4 e6 e24 e-
3
2. Write the symbols for the atoms present in
the structure with the correct configuration.
⇒Generally, the atom with the smallest
electronegativity will be the central atom.
⇒H is always a terminal atom (because it can
form only one bond) and O and the halides
(F, Cl, Br, I) are usually terminal .
O
Cl
C Cl
3. Use lines to indicate electron pair bonds between
pairs of symbols. Use remaining electrons to make
lone pairs so that each atom has an octet (duet for
hydrogen).
O
Cl
Structure and Bonding
C Cl
-
18 e
..
: Cl
..
..
: O:
..
C Cl
.. :
..
: Cl
..
:O :
..
C Cl
.. :
4
4. Indicate the formal charge of the atoms.
⇒ the sum of the formal charges is equal to the
charge of the species.
F.C. = (# val. e-’s) - (# bonds) - (# unshared e -’s)
F.C.(Cl) = 7 - 1 - 6 = 0
F.C.(O) = 6 - 2 - 4 = 0
F.C. (C) = 4 - 4 - 0 = 0
..
: Cl
..
:O :
..
C Cl
.. :
Consider C2 H3 O2 -, acetate
24 e ..
H : O:
H
C
C
H
H
O
C
C
H
..
:O:
H
O
H
:O :
C
C
sp 2
..
:O :
sp3
H
H
H
H
: O:
C
C
..
: O:
H
H
H
..
:O :
C
C
..
O:
H
resonance hybrids
Structure and Bonding
5
Evaluating Resonance Forms
•Resonance forms must have the same structure.
Thus OCN-, NOC -, and CNO- would not be considered as they
have different connectivities.
O
C
N
N
O C
C
N O
The position of the atoms must be the same in all resonance
hybrids.
•Resonance forms in which atoms bonded to one another have
the same charge are unfavorable.
Thus, for the resonance hybrids of FNO2 shown below, hybrid c
is not as important of a contributor as a and b.
. O. :
.. _
. ._
:O
:
: O:
..
:.F.
+
N
a
_
: .O. :
..
:.F.
+
N
b
+
:.F.
: O:
c
+
N
_
: .O. :
•The most important resonance forms of a given resonance
hybrid have the smallest number of formal charges and the
lowest values for these charges. The best forms have no
formal charge at all. Thus, for the cyanate ion, OCN-, c is not
an important contributor to the resonance hybrid.
•The distribution of positive and negative formal charges should be
in agreement with the electronegativities of the atoms.
Consider FNO
Structure and Bonding
6
Exceptions to the Octet Rule
1. If an element has less than four valence electrons,
it may have less than an octet of electrons in a
molecule.
..
:F
..
B
..
..F :
..
: Cl
..
..
Be Cl
.. :
..
: Cl
..
:F:
..
..
Al Cl
.. :
: Cl
.. :
2. Elements from the 3rd period or beyond may have more than
an octet of electrons in a molecule.
PF6- (48 e-’s)
F
F
F
PF
SF4 (34 e-’s)
F
F
F
F
F
S
..
F
The octet rule ⇒ 1 s orbital
and
3 p orbitals
8 electrons in valence shell.
From the 3rd period on, elements can use d orbitals
∴ greater capacity for electrons.
Structure and Bonding
7
POCl3 (32 e-'s)
..
: O:
O
Cl
P
..
: Cl
..
Cl
.. -1
: O:
..
:
Cl
..
P
..
: Cl
..
: Cl
.. :
Cl
: Cl
.. :
.. -1
: O:
..
: Cl
..
: O:
..
: Cl
..
..
:
Cl
..
P
+1
..
:
Cl
..
P
+1
..
:
Cl
..
P
: Cl
.. :
: Cl
.. :
Acids and Bases
Brønsted-Lowry
•Acid--proton donor
•Base--proton acceptor
H
H
O
C
C
O
H
+
O
H
H
acid
base
H3O+ + NH3
Structure and Bonding
H
H
O
C
C
O
+
H
conjugate
base
H2O
+
O
H
H
conjugate
acid
NH4+
8
Acid strength
acid dissociation equilibrium in water
H3O+ + A-
HA + H2O
[H3O + ][A− ]
Ka =
[HA]
pK a = −log K a
Relative acid strengths
Acid
CH 3CH2OH ethanol
16.00
Conjugate Name
base
CH3 CH 2O- ethoxide
H 2O
water
15.74
OH-
hydroxide
HCN
hydrocyanic
acid
acetic acid
9.31
CN-
cyanide
4.76
CH3 CO 2-
acetate
hydrofluoric
acid
nitric acid
3.45
F-
fluoride
-1.3
NO3-
nitrate
-7.0
Cl-
chloride
CH 3CO2H
HF
HNO 3
HCl
Structure and Bonding
Name
hydrochloric
acid
pKa
9
Acids and Bases
Lewis
•Acid--electron pair acceptor
•Base--electron pair donor
Lewis bases have one or more lone pairs of electrons
that can be readily donated:
H
Cl
H +
N
H
-
+
Cl + NH4
H
Lewis
base
Electron deficient compounds are good Lewis Acids
(need an acceptor orbital)
..
:F
..
..
:F
..
..
B ..F :
+
: ..
F:
B
: ..F :
..
: Cl
..
Structure and Bonding
..
F:
..
+
..
Al Cl
.. : +
: Cl
.. :
..
[ :F
.. : ]-
..
:F
..
..
:F:
B
..
F:
..
: ..F :
..
H N H
H
..
[:Cl
.. : ]-
..
: F: H
..
:F B N H
.. +
: F.. : H
..
: Cl
..
..
: Cl :
-
..
Al Cl
.. :
: Cl
.. :
10
Representing molecules
H
•complete (Kekulé) structure
all bonds and atoms are shown
H
C
H
H
H
H
H
C
C
C
C
H
H
H
H
H
•condensed structure
only C-C bonds necessary for clarity are shown
CH3
CH3CH2 CH CH3
CH3
CH3CH2CHCH3
CH3CH2CH(CH3) 2
•skeletal structure
C atoms are not shown
Represented by intersection of lines (bonds) or ends of
terminal lines
H atoms bonded to C are not shown
assume valency is 4 and so we know how many are present
Atoms other than C and H are shown
Structure and Bonding
11
H
Examples,
H
H
H
O
C
C
H
O
O
H
H
OH
H
C
C
C
C
C
H
CH3 COOH
H
H
C
C
H
CH3 C6 H5
H
H
C
H
H
H
C
C
H
H
H
O
C
C
C
C
C
C
H
H
H
C
C
H
H
H
H
C
H
(CH3 )3 CCH2 COC6 H5
Problems
1. Calculate the formal charges on the indicated atoms in each
compound below
..
O:
: Cl
..
P
Cl
. .:
C
O:
:Cl :
2. Draw the Lewis structure of CH2 O and chose the best structure
based on formal charge.
..
:O:
..
H:C:H
Structure and Bonding
:O:
....
H:C:H
lower energy
Lewis structure
12
3. Use δ-/ δ+ convention and the crossed arrow (
) to show
the direction of the expected polarity of the indicated bonds in the
following compounds.
The C-F bond in
fluorobenzene
The C-Si bond in tetramethyl
silane
CH 3
F
H 3C
Si
CH3
CH 3
Structure and Bonding
13