Polar Covalent Bonds: Electronegativity
 Covalent bonds can have ionic character
 These are polar covalent bonds
 Bonding electrons attracted more strongly by one
atom than by the other
 Electron distribution between atoms is not
symmetrical
Bond Polarity and Electronegativity
Symmetrical Covalent Bonds
C–C
C–H
Polar Covalent Bonds
+
C–O
(non-polar)
(polar)
 Electronegativity (EN): intrinsic ability of an atom to
attract the shared electrons in a covalent bond
 Inductive Effect: shifting of sigma bonded electrons
p
to nearby
y electronegative
g
atom
in response
The Periodic Table and Electronegativity
C–H
(
(non-polar)
l )
C - Br and C - I
((polar)
l )
Bond Polarity and Inductive Effect
 Nonpolar
p
Covalent Bonds: atoms with similar EN
 Polar Covalent Bonds: Difference in EN of atoms <
2
 Ionic Bonds: Difference in EN > 2
 C–H bonds, relatively nonpolar C-O, C-X bonds
(more electronegative elements) are polar
 Bonding electrons shift toward electronegative atom
 C acquires
q
p
partial p
positive charge,
g , +
 Electronegative atom acquires partial negative
charge,  Inductive effect: shifting of electrons in a bond in
response to EN of nearby atoms
Electrostatic Potential Maps
 Electrostatic potential
p
maps show calculated
charge distributions
 Colors indicate electronrich (red) and electronpoor (blue)
(bl ) regions
i
 Arrows indicate direction
of bond polarity
Polar Covalent Bonds: Net Dipole
Moments
 Molecules as a whole are often polar from vector summation of
individual bond polarities and lone
lone-pair
pair contributions
 Strongly polar substances soluble in polar solvents like water;
nonpolar substances are insoluble in water.
 Dipole moment () - Net molecular polarity
polarity, due to difference in
summed charges

 - magnitude of charge Q at end of molecular dipole times distance r
between charges

 = Q  r, in debyes (D), 1 D = 3.336  1030 coulomb meter
 length of an average covalent bond, the dipole moment would be 1.60
 1029 Cm, or 4.80 D.
Absence of Net Dipole Moments
 In symmetrical
y
molecules,, the dipole
p
moments of
each bond has one in the opposite direction
 The effects of the local dipoles cancel each other
Drawing Lewis Structures
 Draw molecular skeleton: this will come with practice,




remember atom valence
Determine number of available valence electrons: add an
electron for each negative charge, remove an electron for each
positive charge
Draw all single covalent bonds and lone pairs: give as many
atoms as possible full octets, assigning lone pairs to most
electronegative atoms
Convert lone pairs to multiple bonds if needed: to satisfy
octet rule for as many atoms as possible
Assign formal charges to all atoms: the sum of all charges
must equal total charge of the molecule
Example: Nitromethane: CH3NO2
H
Step 1&2
H
C
O
N
O
H
••
H
H
C
••
O
••
H
C
H
O
N
••
••
O
••
Step 5……………
N
••
H
O
••
••
H
••
Step 3&4
Valence E = 3(1) + 1(4) + 1(5) + 2(6) = 24 e-
••
Formal Charges
 Sometimes it is necessary to have structures with formal charges




on individual
i di id l atoms
t
We compare the bonding of the atom in the molecule to the
valence electron structure
If the atom has one more electron in the molecule, it is shown
with a “-” charge
If the atom has one less electron, it is shown with a “+” charge
Neutral molecules with both a “+” and a “-” are dipolar
 • Atomic sulfur has 6
valence electrons.

Dimethyl suloxide sulfur
has only 5.
 • It has lost an electron and
has positive charge.
 • Oxygen atom in DMSO
has gained electron and has
(-) charge.
Common Formal Charges
charges are typically shown (not necessarily the lone pairs)
add necessary lone pairs and charges to the following compounds:
Resonance
 Some molecules have structures that cannot be shown with a
single
i l representation
i
 In these cases we draw structures that contribute to the final
structure but which differ in the position of the  bond(s) or lone
pair(s)
 Such a structure is delocalized and is represented by
resonance forms
 The resonance forms are connected by a double-headed arrow
We say the pi bond is delocalized, and valence bond theory cannot describe where
Pi bond and one lone pair are located. This is why we must draw multiple versions
And use molecular orbital theory to show delocalization
Curved Arrows and Resonance Forms
 We can imagine
g
that electrons move in p
pairs to
convert from one resonance form to another
 A curved arrow shows that a pair of electrons moves
from the atom or bond at the tail of the arrow to the
atom or bond at the head of the arrow
Drawing Resonance Forms
 Redraw all atoms and sigma
g
bonds in same p
position
 Only move pi electrons and lone pair electrons (note





that electrons must remain with same atom) using
curved arrows to illustrate
All resonance forms must have same net charge
Never violate the octet rule for 2nd row atoms
Resonance forms may be equivalent (degenerate)
but do not have to be
Not all resonance forms contribute equally to the
overall resonance hybrid
There are two additional resonance structures for
nitromethane can you draw them???
nitromethane,
Guidelines to determine major resonance
contributor
ib
 1. Structures with maximum octets are preferred
N
O
N
O
C
major
minor
O
C
O
major
minor
 2. Charges should be located on atoms with compatible
electronegativity
H
H
C
N
N
major
H
C
H
N
N
minor
 3. Structures with minimum charge separation are preferred
O
H
C
O
O
major
H
H
C
O
minor
H
Resonance Hybrids
 A structure with resonance forms does not alternate between
the
h fforms (i
(individual
di id l resonance fforms are iimaginary)
i
)
 Instead, it is a hybrid of all resonance forms, so the structure is
called a resonance hybrid
 For
F example,
l b
benzene (C6H6) h
has ttwo resonance fforms with
ith
alternating double and single bonds
 In the resonance hybrid, the actual structure, all its C-C
bonds equivalent,
equivalent midway between double and single
Different Atoms in Resonance Forms
O
O
-
-
 Are any resonance structures degenerate?
 What is major resonance contributor?
 Can you draw any more resonance structures?
Formal Charge Problems
Provide formal charges for all heteroatoms (non carbon atoms)
atoms).
Remember, all H's on heteroatoms must be (and are) drawn in structure.
H
N
O
H
O
N
H
need to determine number of valence
electrons to know how many lone pairs
to add
For each compound below, provide all lone pairs and hydrogens required to
be consistent with the given formal charges (again hydrogens are provided on
heteroatoms along with any necessary charges).
O
O
H
O
Resonance/Skeletal Problems
O
O
O
O
N N N H
Resonance Practice Problems
1. Draw the important resonance structures for each of the following molecules.
O
H 3C
O
CH2
O
N
H 3C
O
Br
CH3
OCH 3
N
CH3
O
CH3
O
Put in answer keys to the previous slides
also put in slide about hydridization in resonance forms like pyrrole
if sp3 in one, sp2 in another must have sp2 hybrid
y
for p orbital and not move
other atoms
Brønsted Acids and Bases
 A Brønsted acid is a substance that donates
a hydrogen ion (H+)
 A Brønsted
ø
base is a substance that accepts
p
the H+
Acids are shown in red, bases in blue. Curved arrows go from
bases to acids
Ka – the Acidity Constant
 The concentration of water as a solvent does not change
significantly when it is protonated
 The molecular weight of H2O is 18 and one liter weighs 1000
grams, so the concentration is ~ 55.6 M at 25°
 The acidity
y constant,, Ka for HA Ke times 55.6 M ((leaving
g
[water] out of the expression)
 Ka ranges from 1015 for the strongest acids to very small values
(10-60) for the weakest
pKa – the Acid Strength Scale
 p
pKa = -log
g Ka
 The free energy in an equilibrium is related to –log of
Keq (G = -RT log Keq)
 A smaller value of pKa indicates a stronger acid and
is proportional to the energy difference between
products and reactants
 The
Th pK
Ka off water
t is
i 15.74
15 74
Trends in Acidity: What makes a strong
acid?
id
1. St
1
Strength
th off th
the H-A
H Ab
bond
d
The stronger this bond is, the less likely it will break to “give-up” an H+.
For strong acids, this is a relatively weak bond. Important factor when
comparing
p
g bonds between hydrogen
y g and atoms within the same column
(not row) of the periodic table
Why is HI a weaker bond than the others?
pKa =
H-II
H
-10
>
H-Br
H Br
-9
>
H-Cl
H Cl
-7
>
H
H-F
F
3.2
Trends in Acidity: What makes a strong
acid?
id
2. Stability of Conjugate Base,
2
Base AIf the resulting anion from the dissociation is stabilized (A-), it will be easier to
form resulting in a stronger acid. Two major influences operate here:
A: Effect of Electronegativities. If the hydrogen is connected to a very
electronegative atom the bond is polarized and the eventual breaking of the
bond to form A-, a stable anion since A is very electronegative, will become
easier and result in a stronger acid (unless you are comparing atoms in the
same column of the periodic table as in the above example where you must
consider bond strength).
pKa =
H-F
H
F
3.2
>
H-OH
H OH
15.7
>
H-NH
H
NH2
38
>
H-CH
H
CH3
49
Trends in Acidity: What makes a strong
acid?
id
B: Resonance Effects
Effects. Delocalization of charge (through resonance
showing the negative charge is located on more than one atom) will
stabilize the anion, making it easier to form and result in a stronger acid.
O
H3C
C
O
O
H
H3C
C
O
O
H3C
pKa = 5
H3 C
O
H
pKa = 16
H2O versus HNO3
H3 C
O
C
O
Trends in Acidity: What makes a strong
acid?
id
3. Inductive Effects
3
Electronegative groups can inductively “pull” electron density towards
themselves. This can result in a net stabilization of the conjugate base, A-.
This effect is much weaker than that of resonance noted above. An
inductive effect rapidly decreases with distance.
CF3O
O-H
H
pKa =
10
>
CH2FO
FO-H
H
12
>
CH2FCH2O
O-H
H
13
Some Common Organic Acids and their pKa’s
Functional Group
approx pKa
Functional Group
approx pKa
pKa Table: Acidity and Basicity
Conjugate Base
Acid
Strong
Base
CH4
49
CH3
NH3
36
NH2
H C C H
25
H C C
H3C
16
H3 C
Weak
Acid
OH
H2 O
15 7
15.7
H3C SH
11
O
OH
H3 C S
Increasing
Nucleophilicty
NH4
9.2
NH3
H C N
91
9.1
C N
H2S
7.0
SH
H3 C
**
OH
4.8
O
4.7
N N N
H-F
3.2
F
H3O
-1.7
O
S OH
O
-6.5
p-toluene sulfonic acid
Strong
Acid
H3C
H N N N
H3 C
This chart will be VERY
useful for
sub-elim
sub
elim chemistry
O
O
H2O
O
S O
O
H3 C
p-toluene sulfonate or Tosylate
H-Cl
C
-7.0
0
Cl
C
H-Br
-9.0
Br
H-I
-10.0
I
**
Increasing
Leaving Group
Capabilities
Weak
Base
Predicting pKa Values
 Identify most acidic hydrogen and rank from most acidic (1)
to least acidic (4)
(rank by functional groups first, then compare like functional groups)
The reactivity patterns of organic compounds often are acid-base combinations
Predicting Acid–Base Reactions from
pKa Values
l
For any given reaction involving an acid and a base, the equilibrium will be
established such that it favors the weaker acid and weaker base:
H-A
acid
+
B
base
A
+
B-H
conj base
conj acid
A new equilibrium expression can be written for this reaction which simplifies
to the following:
g
Ka reactant acid (HA)
K=
=
10
-(pKa HA - pKa HB)
Ka product acid (HB)
A simple inspection of this new equilibrium expression illustrates that if the reactant acid
has a pKa value 2 units lower than the product conjugate acid then the equilibrium constant
will be 102, or 100. This means that, as the reaction is written, the concentration of products
to reactants will be approximately 100:1. Another way of stating this is that the reaction has
gone to 99% completion (100/101).
Predicting Acid–Base Reactions from
pKa Values
l




Which side of the following
g reaction is favored?
What is K for this reaction?
Does the reaction proceed from left to right?
D
Draw
iin th
the curved
d arrows tto show
h
electron
l t
flflow.
Acids and Bases: The Lewis Definition
 Lewis acids are electron pair acceptors
(have + or + charge,
charge or unfilled octet)
H+, Li+, AlCl3, CH3O-H
 Lewis bases are electron pair donors
(have lone pairs to donate)
H-O-H, CH3-O-H, CH3-NH2
 no scale
l off strengths
t
th as in
i B
Brønsted
t d definition
d fi iti off pK
Ka
Illustration of Curved Arrows in Following
L i Acid-Base
Lewis
A id B
Reactions
R
i
 The combination of a Lewis acid and a Lewis base can be shown with a curved
arrow from
f
base
b
to
t acid
id
Lewis Acids
Lewis Bases
 Lewis bases can accept protons as well as Lewis acids,
therefore the definition encompasses that for Brønsted bases
 Most oxygen- and nitrogen-containing organic compounds are
Lewis bases because they have lone pairs of electrons
 Some compounds
p
can act as both acids and bases,, depending
p
g
on the reaction
C=C -bond will also act as a weak Lewis Base (chapter 6 & 7)
R-X not a Lewis base, but X-1 is
Predicting Reactions
Identifyy Lewis acids,, Lewis bases,, predict
p
product,
p
, add curved arrows
O
N
+
O
H
O
AlCl3
+
O
H
Sample Problems
1 D
1.
Draw ttwo additional
dditi
l resonance structures
t t
with
ith curved
d arrows and
d id
identify
tif
major resonance contributor
2. Predict product if Bronsted-Lowry acid-base reaction occurs. Provide curved arrows.
Which side is fa
favored?
ored?
3. Predict Lewis acid-base product, with curved arrows, for the following reaction.