Bonding and
Molecular Properties
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 in not symmetrical
2
Electronegativity
● Electronegativity is a
measure of the ability
of an atom in a
molecule to attract the
shared electrons in a
covalent bond
Concept proposed by
Linus Pauling
1901-1994
3
Electronegativity
⚫ Atoms on the left of the
periodic table do not
attract electrons as
strongly as nonmetals
on the right side of the
periodic table.
4
Molecular Polarity
Molecules will be polar if
⚫ bonds are polar
AND
⚫ the molecule is NOT “symmetric”
Symmetric molecules
Bond Polarity and Inductive
Effect
⚫ Nonpolar 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 toward electronegative atom
⚫ C acquires partial positive charge, +
⚫ Electronegative atom acquires partial negative charge, -
⚫ Inductive effect: shifting of electrons in a bond in response to
EN of nearby atoms
6
Electrostatic Potential Maps
⚫ Electrostatic
potential maps
show calculated
charge
distributions
⚫ Colors indicate
electron-rich
(red) and
electron-poor
(blue) regions
7
Polar Covalent Bonds: Dipole
Moments
⚫ Molecules as a whole are often polar :from vector summation of
individual bond polarities and lone-pair contributions
⚫ Strongly polar substances soluble in polar solvents like water;
nonpolar substances are insoluble in water.
⚫ Dipole moment - Net molecular 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 Cm, or 4.80 D.
8
Absence of Dipole Moments
⚫ In symmetrical molecules, the dipole moments of each
bond has one in the opposite direction
⚫ The effects of the local dipoles cancel each other
9
Formal Charges
⚫ Sometimes it is necessary to
have structures with formal
charges on individual atoms
⚫ 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
Formal charge = (Group number ) − (Bonding electrons ) − (Number of nonbonding electrons )
1 10
2
Carbon Dioxide, CO2
6 – (1/2)(4) – 4 = 0
O C O
4 – (1/2)(8) – 0 = 0
11
Carbon Dioxide, CO2
6 – (1/2)(2) – 6 = -1
C atom
charge is 0.
O C O
6 – (1/2)(6) – 2 = +1
Which is the predominant resonance structure?
12
Possible Structures
_2 PROBLEMS
.. + + 1) More formal charges (3)
:N.. N O: 2) Multiple (-2) charge
3) Adjacent + charges
_
.. + .. 1) Negative charge is not
:N N O: on the most electroneg-
ative element (oxygen)
+ _
..
:N N O
.. : Looks OK !
13
Rules for Evaluation of Lewis
Diagrams
⚫ Never place like charges on adjacent atoms.
⚫ The total number of charges in the structure should
be kept to a minimum.
⚫ The magnitude of the charges on any atom should
also be kept to a minimum – multiple charges (+2, -
3, etc.) Should be avoided.
⚫ Negative charges should be placed on the most
electronegative elements whenever possible.
⚫ Positive charges should be placed on the least
electronegative elements.
14
Resonance
⚫ Some molecules are have structures that cannot be shown with a
single representation
⚫ 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 to is represented by
resonance forms
⚫ The resonance forms are connected by a double-headed arrow
15
Resonance Hybrids
⚫ A structure with resonance forms does not alternate between the
forms
⚫ Instead, it is a hybrid of the two resonance forms, so the
structure is called a resonance hybrid
⚫ For example, benzene (C6H6) has two resonance forms with
alternating double and single bonds
⚫ In the resonance hybrid, the actual structure, all its C-C
bonds equivalent, midway between double and single
16
Rules for Resonance Forms
⚫ Individual resonance forms are imaginary - the real structure is a
hybrid (only by knowing the contributors can you visualize the
actual structure)
⚫ Resonance forms differ only in the placement of their or
nonbonding electrons
⚫ Different resonance forms of a substance don’t have to be
equivalent
⚫ Resonance forms must be valid Lewis structures: the octet rule
applies
⚫ The resonance hybrid is more stable than any individual
resonance form would be
⚫ The greater the number of resonance forms, the more stable the
substance
17
Curved Arrows and Resonance
Forms
⚫ We can imagine that electrons move in pairs to convert from one
resonance form to another.
⚫ The only difference between resonance forms is in the
distribution of valence electrons.
⚫ The atoms occupy the same place in both resonance forms.
18
Resonance
A curved arrow shows the flow of an electron pair.
The flow may be only:
⚫ From an atom to an adjacent bond or
⚫ From a bond to an adjacent atom or
⚫ From a bond to an adjacent bond.
19
Drawing Resonance Forms
⚫ Any three-atom grouping with a multiple bond
has two resonance forms
20
Different Atoms in Resonance
Forms
⚫ Sometimes resonance forms involve different atom types as well
as locations
⚫ The resulting resonance hybrid has properties associated with
both types of contributors
⚫ The types may contribute unequally
⚫ The “enolate” derived from acetone is a good illustration, with
delocalization between carbon and oxygen
21
Resonance
⚫ Resonance forms must be valid Lewis
structures and obey normal rules of valency.
H H H O
O O
H C C H C C H C C
H O O O
H H
22
Resonance
⚫ The resonance hybrid is
more stable than any
individual resonance
form.
⚫ The greater the number
of resonance forms, the
more stable the
substance.
23
Brønsted Acids and Bases
⚫ “Brønsted-Lowry” is usually shortened to
“Brønsted”
⚫ A Brønsted acid is a substance that donates
a hydrogen ion (H+)
⚫ A Brønsted base is a substance that accepts
the H+
⚫ “proton” is a synonym for H+ - loss of an electron
from H leaving the bare nucleus—a proton
24
Acid-Base Theories
⚫ The Brønsted definition means NH3 is a
BASE in water — and water is itself an ACID
+
H
H N H + H O H H N H + O H
H H
25
Acid-Base Theories
NH3 is a BASE in water — and water is itself an
ACID
+
H
H N H + H O H H N H + O H
H H
NH3 / NH4+ is a conjugate pair — related by
the gain or loss of H+
26
Acid-Base Theories
NH3 is a BASE in water — and water is itself an ACID
+
H
H N H + H O H H N H + O H
H H
NH3 / NH4+ is a conjugate pair — related by the gain
or loss of H+
Every acid has a conjugate base - and vice-versa.
27
Brønsted–Lowry Acids and
Bases
⚫ Conjugate acid-base
pairs differ by a proton,
H+.
H2O/H3O+
NH3/NH4+
NH2- /NH3
28
Weak Acid–Base
⚫ Most acidic substances are weak acids that
only partially ionize in solution.
HA(aq) H+(aq) + A-(aq)
[H3O+] [A-]
Ka =
[HA]
pKa = - log Ka
29
Acid and Base Strength
⚫ The “ability” of a Brønsted acid to donate a proton to
is sometimes referred to as the strength of the acid
(imagine that it is throwing the proton – stronger
acids throw it harder)
⚫ The strength of the acid is measured with respect to
the Brønsted base that receives the proton
⚫ Water is used as a common base for the purpose of
creating a scale of Brønsted acid strength
30
Predicting the Direction of
Acid-Base Reactions
⚫ Based on experiment, we can put acids and bases
on a chart.
ACIDS CONJUGATE BASES
STRONG weak
weak STRONG
⚫ This chart can be used to predict the direction of
reactions between any A-B pair.
⚫ Reactions always go from the stronger A-B pair to
the weaker A-B pair.
31
pKa – the Acid Strength Scale
⚫ pKa = -log 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 pKa of water is 15.74
32
Acid-Base Strengths
⚫ HCN (pKa = 9.31) is ionized in water to a
lesser extent than a HF (pKa = 3.45) solution
of the same concentration.
pKa = -log Ka
33
Predicting Acid–Base Reactions
To deduce on which side the equilibrium lies:
⚫ identify conjugate acid-base pairs (connected above with lines)
⚫ identify the acid on the left side and the acid on the right side of the equation
⚫ the equilibrium lies AWAY FROM the stronger acid.
⚫ The stronger acid has the lower pKa.
O O
C H + O H C + O H
H3C O H3C O
H
pKa = 4.76 pKa = 15.8
An acid with a lower pKa will react with the
conjugate base of an acid with a higher pKa. 34
Predicting Acid–Base
Reactions
⚫ The products must be more stable than the reactants in order for
reaction to occur.
⚫ Therefore, the product acid and base must be weaker and less
reactive than the starting acid and base.
O O
C H + O H C + O H
H3C O H3C O
H
pKa = 4.76 pKa = 15.8
An acid with a lower pKa will react with the
conjugate base of an acid with a higher pKa. 35
Predicting Acid–Base
Reactions
⚫ Ammonia, pKa = 36 and acetone, pKa = 19.
Will the following reaction take place?
O O
+
+ -
+ Na + NH2 + Na + NH3
H3C CH3 H3C CH2
pKa = 36
Weaker acid
pKa = 19
Stronger acid
The reaction will take place as written 36
Predicting Acid–Base Reactions
from pKa Values
⚫ pKa values are related as logarithms to equilibrium
constants
⚫ The difference in two pKa values is the log of the
ratio of equilibrium constants, and can be used to
calculate the extent of transfer
⚫ The weaker base holds the proton more tightly
37
Organic Acids and Organic Bases
⚫ The reaction patterns of organic compounds
often are acid-base combinations
⚫ The transfer of a proton from a strong
Brønsted acid to a Brønsted base, for
example, is a very fast process and will
always occur along with other reactions
38
Organic Acids
⚫ Those that lose a proton from O–H, such as methanol and acetic
acid
⚫ Those that lose a proton from C–H, usually from a carbon atom
next to a C=O double bond (O=C–C–H)
39
Organic Bases
⚫ Have an atom with a lone pair of electrons that can
bond to H+
⚫ Nitrogen-containing compounds derived from
ammonia are the most common organic bases
⚫ Oxygen-containing compounds can react as bases
when with a strong acid or as acids with strong
bases
40
Evaluation of Acid Strength
⚫ In water, all acids form hydronium ion, the important factor of
difference is the conjugate base.
⚫ The difference between a strong acid and a weak acid is in the
stability of the conjugate base.
HA + H2O ¾ H3O+ + A-
A
- WEAK ACID has
strong conj. Base
E (=higher energy)
N
E A-
STRONG ACID has
R ionization weak conj. base
G easier (=lower energy)
Y HA
41
Stabilization Factors
1. Resonance 5. Inductive Effects
2. Electronegativity 6. Charge
3. Size of Atoms 7. Steric Factors
4. Hybridization 8. Solvation Effects
HA + H2O ¾ H3O+ + A-
A
- WEAK ACID has
strong conj. base
E (=higher energy)
N
E A-
STRONG ACID has
R ionization weak conj. base
G easier (=lower energy)
Y HA
42
Resonance Effects
pKa Values
R OH 18 R CH3 45 R NH2 28
OH 10 CH3 30 NH2 25
O O O
R C OH 5 CH3O C CH3 25 R C NH2 15
O
R C CH3 20
O
R C CH2 9
C O
43
R
Resonance in the Acetate Ion
O
CH3 C
O O_
-H+
CH3 C O H
base
acetic acid _
O
CH3 C
O
acetate ion
44
Effect of Electronegativity
increasing
electronegativity pKa Values
O
CH4 RCH3 R C CH3
>45 45 O 20
NH3 RNH2 R C NH2
34 35 O 15
H2O ROH R C OH
16 18 5
HF
3.5
H _ _ _ _
H C H N H O F
H H 45
Acids and Bases: The Lewis
Definition
⚫ Lewis acids are electron pair acceptors and Lewis bases are
electron pair donors
⚫ Brønsted acids are not Lewis acids because they cannot accept
an electron pair directly (only a proton would be a Lewis acid)
⚫ The Lewis definition leads to a general description of many
reaction patterns but there is no scale of strengths as in the
Brønsted definition of pKa
46
Lewis Acids and the Curved Arrow
Formalism
⚫ The Lewis definition of acidity includes metal cations, such as
Mg2+
⚫ They accept a pair of electrons when they form a bond to a
base
⚫ Group 3A elements, such as BF3 and AlCl3, are Lewis acids
because they have unfilled valence orbitals and can accept
electron pairs from Lewis bases
⚫ Transition-metal compounds, such as TiCl 4, FeCl3, ZnCl2, and
SnCl4, are Lewis acids
⚫ Organic compounds that undergo addition reactions with Lewis
bases (discussed later) are called electrophiles and therefore
Lewis Acids
⚫ The combination of a Lewis acid and a Lewis base can shown
with a curved arrow from base to acid
47
Lewis Acids and Bases
48
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 can act as both acids and bases, depending
on the reaction
49
Drawing Chemical Structures
⚫ Chemists use shorthand ways for writing structures
⚫ Condensed structures: Horizontal bonds between C-H and C-C and
single bonds aren't shown but understood
⚫ If C has 3 H’s bonded to it, write CH3
⚫ If C has 2 H’s bonded to it, write CH 2; and so on. The compound called 2-
methylbutane, for example, is written as follows:
⚫ Horizontal bonds between carbons aren't shown in condensed
structures—the CH3, CH2, and CH units are used but vertical bonds are
added for clarity
50
Skeletal Structures
⚫ Minimum amount of information but unambiguous
⚫ C’s not shown, assumed to be at each intersection of two lines
(bonds) and at end of each line
⚫ H’s bonded to C’s aren't shown – whatever number is needed will be
there
⚫ All atoms other than C and H are shown
H
H C H
C C H
H C C
H
H H
51
Drawing Structures
⚫ Condensed structures: carbon–carbon and
carbon–hydrogen bonds not shown.
CH3CH2CH2CH2CH3
⚫ Skeletal Structures
52
Molecular Models
⚫ We often need to visualize the shape or
connections of a molecule in three
dimensions
⚫ Molecular models are three dimensional
objects, on a human scale, that represent
the aspects of interest of the molecule’s
structure (computer models also are
possible)
⚫ Drawings on paper and screens are limited
in what they can present to you
⚫ Framework models (ball-and-stick) are Ball-and-stick
essential for seeing the relationships within
and between molecules – you should own a
set
⚫ Space-filling models are better for
examining the crowding within a molecule
Space-filling
53
Vitamin A, C20H30O
H3C CH3 CH3 H CH3 H H
C CH C C C C
H2C C C C C C OH
H2C C H H H H
CH2 CH3
OH
54
Summary
⚫ Organic molecules often have polar covalent bonds as a result of
unsymmetrical electron sharing caused by differences in the electronegativity
of atoms
⚫ The polarity of a molecule is measured by its dipole moment, .
⚫ (+) and (−) indicate formal charges on atoms in molecules to keep track of
valence electrons around an atom
⚫ Some substances must be shown as a resonance hybrid of two or more
resonance forms that differ by the location of electrons.
⚫ A Brønsted(–Lowry) acid donates a proton
⚫ A Brønsted(–Lowry) base accepts a proton
⚫ The strength Brønsted acid is related to the -1 times the logarithm of the acidity
constant, pKa. Weaker acids have higher pKa’s
⚫ A Lewis acid has an empty orbital that can accept an electron pair
⚫ A Lewis base can donate an unshared electron pair
⚫ In condensed structures C-C and C-H are implied
⚫ Skeletal structures show bonds and not C or H (C is shown as a junction of
two lines) – other atoms are shown
⚫ Molecular models are useful for representing structures for study
55
