Chapter 2
Introduction
• The Lewis Model of Bonding:
Atoms bond together in such a way that each
atom participating in a chemical bond acquires a
completed valence-shell electron configuration
resembling that of the noble gas nearest it in the
Periodic Table.
• Ionic Bond – is a chemical bond resulting from
the electrostatic attraction between an anion and
a cation.
Na
+
..
F
..
Na+ +
..
F
..
• Covalent Bond – is a chemical bond formed by the
sharing of electron pairs between atoms.
2 H·
H2
• Based on the the degree of electron sharing, covalent
bonds can be divided into:
• Nonpolar covalent bond
• Polar covalent bond
I. Polar Covalent Bonds
A.
Electronegativity
B.
Dipole Moments
• Polar Covalent Bond – is a covalent bond with
ionic character.
 Bonding electrons are attracted more strongly by one
atom than by the other
 Electron distribution between atoms is not symmetrical
A. Polar Covalent Bonds: Electronegativity
• Electronegativity (EN):
 is a measure of the force of an atom’s attraction
for electrons it shares in a covalent bond with
another atom.
 is an intrinsic ability of an atom to attract the
shared electrons in a covalent bond
Bond Polarity and Electronegativity
• EN of an atom is related to its ionization energy
(IE) and electron affinity (EA).
• Atom in ground state (g) + Energy  Atom+ (g) + e- DE = IE
• Atom (g) + e-  AtomDE = EA
• Differences in EN produce bond polarity.
Bond Polarity
~
DEN
~
DEN
SEN
Bond Polarity and Electronegativity
• EN is based on an arbitrary scale.
• The most widely used scale of EN was devised by
Linus Pauling in the 1930’s and is based on bond
energies.
• Important EN values:
• F is the most electronegative (EN = 4.0)
• Cs is the least electronegative (EN = 0.7)
• C has an EN = 2.5
The Periodic Table and Electronegativity
The Periodic Table and Electronegativity
• Metals on left side of periodic table attract electrons
weakly, lower EN
The Periodic Table and Electronegativity
• Halogens and other reactive nonmetals on right side
of periodic table attract electrons strongly, higher EN
Classification of Covalent Bonds
Difference in EN
between bonded Atoms
Type of Bond
Similar EN
Nonpolar Covalent
Less than 2
Polar Covalent
Greater than 2
Ionic
Classification of Covalent Bonds
Difference in EN
between bonded Atoms
Type of Bond
Less than 0.5
Nonpolar Covalent
0.5 to 1.9
Polar Covalent
Greater than 1.9
Ionic
Classification of Covalent Bonds: Examples
• C–H bonds are relatively nonpolar
DEN = ENC – ENH = 2.5 - 2.1 = 0.4
• C-O and C-X bonds (more electronegative
elements) are polar
DEN = ENO – ENC = 3.5 - 2.5 = 1
Bond Polarity and Inductive Effect
• Bonding electrons are drawn toward electronegative
atom


C acquires partial positive charge, +
Electronegative atom acquires partial negative
charge, -
The crossed arrow indicates
the direction of the electron
displacement
Bond Polarity and Inductive Effect
• Inductive Effect – is the shifting of electrons in a
bond in response to EN of nearby atoms
 Metals (Li and Mg) inductively donate e-
 Reactive nonmetals (O and Cl) inductively withdraw e-
Electrostatic Potential Maps
• Electrostatic
potential maps
show calculated
charge distributions
• Colors indicate
electron-rich (red)
and electron-poor
(blue) regions
Practice Problem: Which element in each of the following pairs
is more electronegative?
a. Li or H
b. B or Br
c. Cl or I
d. C or H
Practice Problem: Use the +/- convention to indicate the
direction of expected polarity for each of the
bonds indicated
a. H3C — Br
e. H3C — OH
b. H3C — NH2
f. H3C — MgBr
c. H3C — Li
g. H3C — F
d. H2N — H
Practice Problem: Use the electronegativity values to rank the
following bonds from least polar to most polar:
H3C — Li
H3C — K
H3C — F
H3C — MgBr
H3C — OH
Practice Problem: Look at the following electrostatic potential
map of methyl alcohol, and tell the direction of
polarization of the C-O bond:
B. Polar Covalent Bonds: Dipole Moments
• Molecular Polarity
 is the tendency of molecules as a whole to be polar
 results from vector summation of individual bond
polarities and lone-pair contributions
• “Like dissolves like”
• Strongly polar substances are soluble in polar
solvents like water; nonpolar substances are
insoluble in water.
•
To predict whether a molecule is polar,
determine:
1. if the molecule has polar bonds, and
2. the arrangements of these bonds in space
• Dipole moment (m) – is a measure of net
molecular polarity, due to difference in summed
charges
m= Q  r
magnitude of charge Q at
either end of molecular dipole
distance r between charges
• It is expressed in Debyes (D)
• 1 D = 3.336 x 10-30 Coulomb meter
Calculating the dipole moment of an average bond:
raverage covalent bond = 100pm
Qelectron = 1.60 x 10-19 C
The dipole moment (m) of an average covalent
bond is 4.80 D
Calculating Ionic Character: Chloromethane
Given that r C-Cl = 178 pm and assuming mC-H is negligible, then
m calculated CH3Cl = 178 pm x 4.80 D = 8.5 D
mmeasured CH3Cl = 1.87D
% ionic character = mmeasured /
mcalculated x 100 = 22%
Dipole Moments in Water and Ammonia
• H2O and NH3 have large dipole moments:
 ENO and ENN > ENH
 Both O and N have lone-pair electrons oriented away
from all nuclei
Dipole Moments in Water and Ammonia
• H2O and NH3 have large dipole moments:
 ENO and ENN > ENH
 Both O and N have lone-pair electrons oriented away
from all nuclei
Absence of Dipole Moments
• In symmetrical molecules, the dipole moment of each
bond has one in the opposite direction
• The effects of the local dipoles cancel each other
Practice Problem: Carbon dioxide, CO2, has zero dipole moment
even though carbon-oxygen bonds are
strongly polarized. Explain.
Practice Problem: Make three-dimensional drawings of the
following molecules, and predict whether
each has a dipole moment. If you expect a
dipole moment, show its direction.
a. H2C = CH2
b. CHCl3
c. CH2Cl2
d. H2C = CCl2
II. Formal Charges and Resonance
A.
Formal Charges
B.
Resonance
A. Formal Charges
• Formal Charge - is the charge on an atom in
a molecule or polyatomic ion
Comparing the bonding of the atom in the molecule to
the valence electron structure
No formal charge
No formal charge
Calculating the formal charge
• 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
Practice Problem: Dimethyl sulfoxide, a common solvent, has
the structure indicated. Show why dimethyl
sulfoxide must have formal charges on S and
O.
Practice Problem: Calculate formal charges for the nonhydrogen
atoms in the following molecules:
a. Diazomethane,
b. Acetonitrile oxide,
c. Methyl isocyanide
Practice Problem: Organic phosphates occur commonly among
biological molecules. Calculate formal
charges on the four O atoms in the methyl
phosphate ion.
B. Resonance
• Some molecules have Lewis 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)
Resonance Forms
• Resonance forms – are Lewis structures of the same
molecule whose only difference is the placement of  and
nonbonding valence electrons (= delocalized).
• The atoms occupy the same place in the different forms
• The connections between atoms are the same.
• The resonance forms are connected by a double-headed
arrow
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
Resonance Hybrids: Benzene
• 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
are equivalent, midway between double and single
Rules for Resonance Forms1
1. Individual resonance forms are imaginary - the real
structure is a hybrid (only by knowing the contributors
can you visualize the actual structure)
Rules for Resonance Forms2
2.
Resonance forms differ only in the placement of their
 or nonbonding electrons
Rules for Resonance Forms3
3.
Different resonance forms of a substance don’t have
to be equivalent
Rules for Resonance Forms4
4.
Resonance forms must be valid Lewis structures: the
octet rule applies
Rules for Resonance Forms5
5.
The resonance hybrid is more stable than any
individual resonance form
•
Resonance leads to stability. The larger the # of the
resonance forms, the more stable the substance
Rules for Resonance Forms
1. Individual resonance forms are imaginary - the real
structure is a hybrid (only by knowing the contributors can
you visualize the actual structure)
2. Resonance forms differ only in the placement of their  or
nonbonding electrons
3. Different resonance forms of a substance don’t have to be
equivalent
4. Resonance forms must be valid Lewis structures: the octet
rule applies
5. The resonance hybrid is more stable than any individual
resonance form
Curved Arrows and Resonance Forms
•
We can imagine that electrons move in 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
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
Drawing Resonance Forms
•
Any three-atom grouping with a p orbital on each
atom has two resonance forms
X, Y, Z can be C, N, O, P or S
Resonance Structures: 2,4-Pentanedione
• The anion derived from 2,4-pentanedione
– Lone pair of electrons and a formal negative
charge on the central carbon atom, next to a C=O
bond on the left and on the right
– Three resonance structures result
Relative Importance of contributing Structures
The most important contributing structures have:
1.
2.
3.
4.
filled valence shells
a maximum number of covalent bonds
the least separation of unlike charge, and
any negative charge on a more electronegative atom
and/or any positive charge on a less electronegative
atom
Practice Problem: Draw the indicated number of resonance
structures for each of the following species:
a. The nitrate ion, NO3- (3)
b. The allyl cation, H2C=CH-CH2+ (2)
c. Hydrazoic acid,
d. (2)
(2)
III. Acids and Bases1
A.
The Brønsted-Lowry Definition
B.
Acid and Base Strength
C.
Predicting Acid-Base Reactions
from pKa Values
III. Acids and Bases2
D.
Organic Acids and Organic Bases
E.
The Lewis Definition
Introduction
 The terms “acid” and “base” can have different
meanings in different contexts.
• For that reason, we specify the usage with more
complete terminology
• There are three definitions of acid-base: Arrhenius,
Bronsted-Lowry, and Lewis
• According to the Arrhenius definition:
• An Arrhenius acid - is a substance that donates H+ in
aqueous solutions
• An Arrhenius base - is a substance that donates OHin aqueous solutions
• The Arrhenius definition is not useful in organic
chemistry.
A. The Brønsted-Lowry Definition
• Brønsted–Lowry theory defines acids and
bases by their role in reactions that transfer
protons (H+) between donors and acceptors
• “Brønsted-Lowry” is usually shortened to
“Brønsted”
Brønsted Acids and Bases
• 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+ (i.e. Loss of the
valence electron from a neutral H leaves only the
hydrogen nucleus - a proton)
Brønsted acid–base Reactions
The Reaction of HCl with H2O
• When HCl gas dissolves in water, a Brønsted
acid–base reaction occurs
• HCl donates a proton to water molecule,
yielding hydronium ion (H3O+) and Cl-
Other Brønsted acid–base Reactions
• The reverse is also a Brønsted acid–base
reaction of the conjugate acid and conjugate
base
Practice Problem: Nitric acid (HNO3) reacts with ammonia (NH3)
to yield ammonium nitrate. Write the reaction,
and identify the acid, the base, the conjugate
acid product, and the conjugate base product.
B. Acid and Base Strength
• The equilibrium constant (Keq) for the reaction
of an acid (HA) with water to form hydronium
ion and the conjugate base (A-) is a measure
related to the strength of the acid
• Stronger acids have larger Keq
Ka – the acidity constant
• The concentration of water as a solvent does
not change significantly when it is protonated
 Molecular weight (H2O) = 18 g/mol.
 Density (H2O) = 1000 g/l
 [H2O] ~ 55.6 M at 25°C
Ka – the acidity constant
• The acidity constant, Ka for HA is
• Ka ranges from 1015 for the strongest acids to
very small values (10-60) for the weakest
Acid-Base Strength
• The “ability” of a Brønsted acid to donate a
proton is sometimes referred to as the
strength of the acid.
• 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
pKa – The acid strength scale
• Acid strength is expressed using pKa values:
pKa = - log Ka
• The free energy in an equilibrium is related to log of Keq:
DG = -RT ln Keq
• A larger value of Ka indicates a stronger acid
and is proportional to the energy difference
between products and reactants
pKa – The acid strength scale
• The pKa of water is 15.74
• The stronger the acid, the weaker its conjugate base.
The weaker the acid, the stronger the conjugate
base.
Practice Problem: Formic acid, HCO2H, has pKa = 3.75, and
picric acid, C6H3N3O7, has pKa = 0.38. Which
is the stronger acid?
Practice Problem: Amide ion, H2N-, is a much stronger base
than hydroxide ion, HO-. Which would you
expect to be a stronger acid, NH3 or H2O?
Explain.
C. 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:
DpKa = log Keq2/Keq1
 DpKa can be used to calculate the extent of
H+ transfer
• H+ will always go from the stronger acid to the
stronger base
• The product acid must be weaker and less
reactive than the starting acid
• The product base must be weaker and less
reactive than the starting base
• The stronger base holds the proton more
tightly
Practice Problem: Will either of the following reactions take
place as written, according to the pKa data?
?
a. HCN + CH3CO2- Na+  Na+ -CN + CH3CO2H
?
b. CH3CH2OH + Na+-CN  CH3CH2O-Na+ + HCN
Practice Problem: Ammonia, NH3, has pKa = 36 and acetone
has pKa = 19. Will the following reaction take
place?
Practice Problem: What is the Ka of HCN if its pKa = 9.31?
D.
Organic Acids and Organic Bases
• The reaction patterns of organic compounds
often are acid-base combinations
• The transfer of a H+ from a strong Brønsted
acid to a Brønsted base
– is a very fast process
– will always occur along with other reactions
Organic Acids
• Organic acids lose a proton either from:
• O–H, such as methanol and acetic acid
• C–H, usually from a carbon atom next to a
C=O double bond (O=C–C–H)
Organic Acids
• The conjugate base resulting from loss of H+ is
stabilized by having its negative charge on a highly
electronegative oxygen atom  Acidity
Organic Acids
Organic Bases
• 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
Organic Bases
E.
The Lewis Definition
• A Lewis acid is an electron pair acceptor; it
has a low-energy empty orbital
• A Lewis base is an electron pair donor
• 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
Lewis Acids
Lewis acids include:
• H+
• 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: They
have unfilled valence orbitals and can accept electron
pairs from Lewis bases
• Transition-metal compounds, such as TiCl4, FeCl3,
ZnCl2, and SnCl4: They have unfilled valence orbitals
and can accept electron pairs from Lewis bases
Lewis Acids
Lewis Acids
• Organic compounds that undergo addition
reactions with Lewis bases are called
electrophiles and therefore Lewis Acids
• The combination of a Lewis acid and a Lewis
base can be shown with a curved arrow from
base to acid
Illustration of Curved Arrows in Following Lewis
Acid-Base Reactions
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
Lewis Bases
Practice Problem: Using curved arrows, show how the species
in part (a) can act as Lewis bases in their
reactions with HCl, and show how the species
in part (b) can act as Lewis acids in their
reaction with OH-
a. CH3CH2OH, HN(CH3)2, P(CH3)3
b. H3C+, B(CH3)3, MgBr2
Practice Problem: Explain by calculating formal charges why the
following acid-base reaction products have
the charges indicated:
+
a. F3B — O — CH3
|
CH3
CH3
|
+
b. Cl3Al — N — CH3
|
CH3
Practice Problem: The organic compound imidazole can act as
both an acid and a base. Look at the following
electrostatic potential map, and identify the
most acidic hydrogen atom and the most
basic nitrogen in imidazole
IV. Chemical Structures
A.
Drawing Chemical Structures
B.
Molecular Models
A.
Drawing Chemical Structure
• Chemists use shorthand ways for writing structures.
• Organic molecules are usually drawn as:
• Condensed structures
• Skeletal structures
Condensed Structures
• C-H and C-C 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 CH2; and so on.
• Horizontal bonds between carbons aren't shown in
condensed structures—the CH3, CH2, and CH units
are simply placed next each other but vertical bonds
are added for clarity
Skeletal Structures
• Minimum amount of information but
unambiguous
• Rules for Skeletal Structures:
1. C’s are not shown, assumed to be at each
intersection of two lines (bonds) and at end of
each line
2. H’s bonded to C’s aren't shown – whatever
number is needed will be there
3. All atoms other than C and H are shown
Practice Problem: Tell how many hydrogens are bonded to each
carbon in the following compounds, and give
the molecular formula of each substance:
Practice Problem: Propose skeletal structures for compounds
that satisfy the following molecular formulas
(there is more than one possibility in each
case):
a. C5H12
b. C2H7N
c. C3H6O
d. C4H9Cl
Practice Problem: The following molecular model is a
representation of para-aminobenzoic acid
(PABA), the active ingredient in many
sunscreens. Indicate the positions of the
multiple bonds, and draw a skeletal structure
(gray = C, red = O, blue = N, ivory = H)
B. Molecular Model
• 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)
Framework
Space-filling
• Drawings on paper and
screens are limited in what
they can present to you
• Framework models (ball-andstick) are 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
Framework
Space-filling
Chapter 2