Transcript 2. Polar Covalent Bonds: Acids and Bases Why this chapter?

2. Polar Covalent Bonds:
Acids and Bases
Why this chapter?
Description of basic ways chemists account for
chemical reactivity.
Establish foundation for understanding specific
reactions discussed in subsequent chapters.
2.1 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
2
Bond Polarity and Electronegativity







Electronegativity (EN): intrinsic ability of an atom to
attract the shared electrons in a covalent bond
Differences in EN produce bond polarity
Arbitrary scale. As shown in Figure 2.2,
electronegativities are based on an arbitrary scale
F is most electronegative (EN = 4.0), Cs is least (EN =
0.7)
Metals on left side of periodic table attract electrons
weakly, lower EN
Halogens and other reactive nonmetals on right side of
periodic table attract electrons strongly, higher
electronegativities
EN of C = 2.5
3
The Periodic Table and Electronegativity
4
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
5
Polarity, Shape
CO2
O C
3.5
2.5
O
3.5
O=C=O
-
+ -
Polar Covalent Bonds
Linear Shape (180o)
Nonpolar Compound
Electronegativity
Difference
< 0.5 Nonpolar
0.5-1.7 Polar
>1.8 Ionic
Polarity, Shape
e-’s in 2 directions = 180o
O=C=O
-
Linear
+ Nonpolar Compound
e-’s in 3 directions = 120o
-O
Trigonal planar
C
Polar Compound
H + H
Polarity, Shape
BF3
F
F B F
F
-
2.0
-
4.0
B +
4.0 F
F
4.0
(120o)
Trigonal Planar
Polar Covalent Bonds
Nonpolar Compound
Polarity, Shape
H
e-’s in 4 directions = 109.5o
H
H
H C H H C H
Cl
Cl
4 directions = 109.5o
H O H
H-O-H
+
+
HH
Bent
-
Cl
Tetrahedral
H
O
C
H+
Polarity, Shape
e-’s in 4 directions = 109.5o
-
H N H H N H +H
H
+
H
H
N
H+
Pyramidal
(109.5o)
Tetrahedral Configuration of Electrons
Trigonal Pyramid Configuration of Atoms
Tetrahedral electron-pair
Geometries
Tetrahedral
Pyramidal
Bent
Electrostatic Potential Maps
Electrostatic potential
maps show calculated
charge distributions
 Colors indicate electronrich (red) and electronpoor (blue) regions
 Arrows indicate direction
of bond polarity

15
2.2 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
Cm, or 4.80 D.
16
Dipole Moments in Water and Ammonia

Large dipole moments
◦ EN of O and N > H
◦ Both O and N have lone-pair electrons oriented away from
all nuclei
17
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
18
Learning Check:
H
F
F
F
F
H
F
F
F
H
H
H
F
H
F
F
A
B
C
D
Which of the above molecules is the most polar?
1.
2.
3.
4.
5.
A
B
C
D
It cannot be
determined
Solution:
H
F
F
F
F
H
F
F
F
H
H
H
F
H
F
F
A
B
C
D
Which of the above molecules is the most polar?
1.
2.
3.
4.
5.
A
B
C
D
It cannot be
determined
Learning Check:
Which of the following molecules is nonpolar?
1.
2.
3.
4.
5.
H2O
CH3CN
CHCl3
CH2Cl2
CO2
21
Solution:
Which of the following molecules is nonpolar?
1.
2.
3.
4.
5.
H2O
CH3CN
CHCl3
CH2Cl2
CO2
22
Learning Check:
H
F
F
H
C
H
F
A
C C
F
H
B
H
F
H
F
C
Which of the above molecules has (have) a dipole
moment?
1.
2.
3.
4.
5.
A only
B only
C only
A and B
B and C
Solution:
H
F
F
H
C
H
F
A
C C
F
H
B
H
F
H
F
C
Which of the above molecules has (have) a dipole
moment?
1.
2.
3.
4.
5.
A only
B only
C only
A and B
B and C
2.3 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
25
Formal Charge for Dimethyl Sulfoxide
• 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.
26
27
Learning Check:
What are the hybridization and a formal charge,
respectively, on boron in BH4–?
1.
2.
3.
4.
5.
sp3, 0
sp3, –1
sp3, +1
sp2, –1
sp2, +1
Solution:
What are the hybridization and a formal charge,
respectively, on boron in BH4–?
1.
2.
3.
4.
5.
sp3, 0
sp3, –1
sp3, +1
sp2, –1
sp2, +1
2.4 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 is represented by resonance
forms
The resonance forms are connected by a double-headed arrow
30
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 are
equivalent, midway between double and single
31
2.5 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

32
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

33
2.6 Drawing Resonance Forms

Any three-atom grouping with a multiple
bond has two resonance forms
34
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
35
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
36
Learning Check:
Which of the following statements is false
concerning resonance structures?
1.
2.
3.
4.
5.
The arrangement of nuclei in all contributing structures
must be the same.
The arrangement of electrons in each contributing
structure is different.
Each atom in a contributing structure must have a
completed valence shell.
The contributing structures may have different energies.
All contributing structures must have the correct number
of valence electrons.
Solution:
Which of the following statements is false
concerning resonance structures?
1.
2.
3.
4.
5.
The arrangement of nuclei in all contributing structures
must be the same.
The arrangement of electrons in each contributing
structure is different.
Each atom in a contributing structure must have a
completed valence shell.
The contributing structures may have different energies.
All contributing structures must have the correct number
of valence electrons.
Learning Check:
Consider all equivalent resonance structures for the carbonate
anion shown below. If the bond order is defined as the number of
electron pairs shared between two atoms, what is the average
bond order for the C-O1 bond?
1
1.
2.
3.
4.
5.
2
3/2
4/3
1
1/4
O
O
O
Solution:
Consider all equivalent resonance structures for the carbonate
anion shown below. If the bond order is defined as the number of
electron pairs shared between two atoms, what is the average
bond order for the C-O1 bond?
1
1.
2.
3.
4.
5.
2
3/2
4/3
1
1/4
O
O
O
Learning Check:
Consider all equivalent resonance structures for
carbonate anion shown below. What is the formal
charge on O1?
1
O
1.
2.
3.
4.
5.
–1
–2/3
–1/2
–1/3
0
O
O
Solution:
Consider all equivalent resonance structures for
carbonate anion shown below. What is the formal
charge on O1?
1
O
1.
2.
3.
4.
5.
–1
–2/3
–1/2
–1/3
0
O
O
Learning Check:
Anions with the negative charge residing on carbons are
called carbanions. In the methyl anion the carbon is sp3
hybridized. What is the hybridization of carbon 1 in the
allyl anion?
H
1 H
H
H
H
1.
sp3
2.
sp2
sp
sp3d
carbon 1 is not hybridized
3.
4.
5.
allyl anion (a carbanion)
Solution:
Anions with the negative charge residing on carbons are
called carbanions. In the methyl anion the carbon is sp3
hybridized. What is the hybridization of carbon 1 in the
allyl anion?
H
1 H
H
H
H
1.
sp3
2.
sp2
sp
sp3d
carbon 1 is not hybridized
3.
4.
5.
allyl anion (a carbanion)
Learning Check:
How many uncharged resonance structures
exist for the following molecule?
1.
2.
3.
4.
5.
Only the one shown above
2
3
4
5
Solution:
How many uncharged resonance structures
exist for the following molecule?
1.
2.
3.
4.
5.
Only the one shown above
2
3
4
5
2.7 Acids and Bases: The Brønsted–
Lowry Definition




The terms “acid” and “base” can have different meanings
in different contexts
For that reason, we specify the usage with more complete
terminology
The idea that acids are solutions containing a lot of “H+”
and bases are solutions containing a lot of “OH-” is not
very useful in organic chemistry
Instead, Brønsted–Lowry theory defines acids and bases
by their role in reactions that transfer protons (H+)
between donors and acceptors
47
Bronsted-Lowry theory
Acid =
a Proton donor
donates a proton, H+, to solvent
Example: HCl
HCl + H2O
H3O+ + Clhydronium ion
Bronsted-Lowry theory
Example: HNO3
HNO3 + H2O
HNO3 is the acid:
H3O+ + NO3it donates the proton.
Bronsted-Lowry theory
Base =
a proton acceptor
accepts a Proton, H+, from solvent
NH3 + H2O
NH4+ + OH-
Accepted H+
NH3 is the Base:
it takes the proton.
Hydroxide ions (OH-) form when water
donates the proton.
Common Acids


Battery Acid
Stomach Acid
H2SO4
HCl
Sulfuric Acid
Hydrochloric Acid

Coca Cola
H3PO4
Phosphoric Acid


Carbonated Water H2CO3
Vinegar
HC2H3O2
Citrus fruits
H3C6O7H8
Vitamin C
HC6O6H7
H2C4O6H4
Grapes
H2C9O4H8
Aspirin




Carbonic Acid
Acetic Acid
Citric Acid
Ascorbic Acid
Tartaric Acid
Acetyl Salicylic Acid
Naming Acids
•Binary Acids = (-ide ending)
Hydro- ______-ic
Acid
HCl = Hydrogen Chloride = Hydrochloric
Acid
•Oxoacids = H, Nonmetal & O
-ate ending (The higher number of O’s)
______________-ic
Acid
HNO3 = Hydrogen Nitrate = Nitric Acid
-ite
ending (The lower number of O’s)
______________-ous Acid
HNO2 = Hydrogen Nitrite = Nitrous Acid
Naming Acids
Non-Acid name
H2SO4
H2SO3
H2S
H3PO4
HBr
H2CO3
Dihydrogen Sulfate
Acid Name
Sulfuric Acid
Dihydrogen Sulfite
Sulfurous Acid
Dihydrogen Sulfide
Hydrosulfuric Acid
Trihydrogen Phosphate
Hydrogen Bromide
HNO2
Dihydrogen carbonate
Hydrogen Nitrite
HNO3
Hydrogen Nitrate
HC2H3O2 Hydrogen Acetate
Phosphoric Acid
Hydrobromic Acid
Carbonic Acid
Nitrous Acid
Nitric Acid
Acetic Acid
Naming Acids
Non-Acid name
H2SO4
H2SO3
H2SO2
HCl
HClO
HClO2
Dihydrogen Sulfate
Dihydrogen Sulfite
Dihydrogen Sulfide
Hydrogen Chloride
Hydrogen hypochlorite
Hydrogen chlorite
HClO3 Hydrogen chlorate
HClO4
Hydrogen Perchlorate
Acid Name
Sulfuric Acid
Sulfurous Acid
Hyposulfurous Acid
Hydrochloric Acid
Hypochlorous Acid
Chlorous Acid
Chloric Acid
Perchloric Acid
Battery Acid
Stomach Acid
H2SO4
HCl
Coca Cola
Carbonated Water
Vinegar
Citrus fruits
Vitamin C
Grapes
Aspirin
H3PO4
H2CO3
HC2H3O2
H3C6O7H8
HC6O6H7
H2C4O6H4
H2C9O4H8
Common Acids
Strong
•100% ionization
•Strong electrolyte
Weak
•Partial ionization
•Weak electrolyte
•Taste sour
Acids
H-Cl + H-O-H
H
H
O+
H-Cl
H
H
O+
H
+
Cl-
HC2H3O2
Cl-
HC2H3O2
H+
H
C2H3O21-
H
H-C2H3O2 + H-O-H
+
+
C
H
O
2 3 2
O
H
H
Bases
Base = gives hydroxide ions in water.
(Arrhenius definition)
= takes hydrogen ions in water.
(Bronsted-Lowry definition)
NaOH
NaOH
Na+ + OH-
Na+
OH-
OHNa+
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
58
The Reaction of Acid with Base



Hydronium ion, product when base H2O gains a proton
HCl donates a proton to water molecule, yielding hydronium ion
(H3O+) [conjugate acid] and Cl [conjugate base]
The reverse is also a Brønsted acid–base reaction of the conjugate
acid and conjugate base
59
2.8 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
 Note that brackets [ ] indicate concentration, moles per
liter, M.

60
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 ~ 55.4 M at 25°
The acidity constant, Ka for HA Keq times 55.6 M (leaving [water]
out of the expression)
Ka ranges from 1015 for the strongest acids to very small values (10-60)
for the weakest
61
62
pKa – the Acid Strength Scale




pKa = -log Ka
The free energy in an equilibrium is related to –log of Keq
(DG = -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
63
2.9 Predicting Acid–Base Reactions
from pKa Values




pKa values are related as logarithms to equilibrium constants
Useful for predicting whether a given acid-base reaction will take
place
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 stronger base holds the proton more tightly
64
2.10 Organic Acids and Organic Bases

-
Organic Acids:
characterized by the presence of
positively polarized hydrogen atom
65
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)

66
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

67
2.11 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

68
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 TiCl4, 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
69
Illustration of Curved Arrows in Following
Lewis Acid-Base Reactions
70
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
71
Learning Check:
All Lewis acids can also be classified as
Brønsted-Lowry acids.
1.
2.
True
False
Solution:
All Lewis acids can also be classified as
Brønsted-Lowry acids.
1.
2.
True
False
Learning Check:
Which statement best describes the following
reaction?
NH3 +
1.
2.
3.
4.
5.
BH3
H3N BH3
an acid-base reaction where NH3 acts as a BrønstedLowry acid
an acid-base reaction where NH3 acts as a BrønstedLowry base
an acid-base reaction where NH3 acts as a Lewis acid
an acid-base reaction where NH3 acts as a Lewis base
None of these adequately describes the reaction.
Solution:
Which statement best describes the following
reaction?
NH3 +
1.
2.
3.
4.
5.
BH3
H3N BH3
an acid-base reaction where NH3 acts as a BrønstedLowry acid
an acid-base reaction where NH3 acts as a BrønstedLowry base
an acid-base reaction where NH3 acts as a Lewis acid
an acid-base reaction where NH3 acts as a Lewis base
None of these adequately describes the reaction.
Learning Check:
What is the main reason why molecule A is less
acidic than molecule B?
A: CH3CH2CH2-H
B: CH3CH2O-H
1.
2.
3.
4.
5.
There are more hydrogens in A.
The C-H bond is weaker than the O-H bond.
The conjugate base of A is more stable than the conjugate
base of B.
In the conjugate base of B, the negative charge resides on a
more electronegative atom.
There is resonance and inductive stabilization of conjugate
base of B.
Solution:
What is the main reason why molecule A is less
acidic than molecule B?
A: CH3CH2CH2-H
B: CH3CH2O-H
1.
2.
3.
4.
5.
There are more hydrogens in A.
The C-H bond is weaker than the O-H bond.
The conjugate base of A is more stable than the conjugate
base of B.
In the conjugate base of B, the negative charge resides on a
more electronegative atom.
There is resonance and inductive stabilization of conjugate
base of B.
Learning Check:
Ammonia can act as all four classifications of acids or
bases (Lewis acid, Lewis base, Brønsted-Lowry acid,
and Brønsted-Lowry base.)
1.
2.
True
False
Solution:
Ammonia can act as all four classifications of acids or
bases (Lewis acid, Lewis base, Brønsted-Lowry acid,
and Brønsted-Lowry base.)
1.
2.
True
False
Learning Check:
Van der Waals interactions refer to the weak attractive forces
that occur between individual molecules and are very
important in determining the shape and properties of
biological molecules such as proteins. Which of the following
play no role in van der Waals interactions?
1.
2.
3.
4.
5.
hydrogen bonding
dispersion forces
dipole-dipole forces
magnetic forces
All of these are important in van der Waals
interactions.
Solution:
Van der Waals interactions refer to the weak attractive forces
that occur between individual molecules and are very
important in determining the shape and properties of
biological molecules such as proteins. Which of the following
play no role in van der Waals interactions?
1.
2.
3.
4.
5.
hydrogen bonding
dispersion forces
dipole-dipole forces
magnetic forces
All of these are important in van der Waals
interactions.
Learning Check:
Which is the correct order of the pKa
values?
1.
2.
3.
4.
5.
H2O > HO– > H3O+
HO– > H3O+ > H2O
H3O+ > HO – > H2O
HO– > H2O > H3O+
H3O+ > H2O > HO–
Solution:
Which is the correct order of the pKa
values?
1.
2.
3.
4.
5.
H2O > HO– > H3O+
HO– > H3O+ > H2O
H3O+ > HO – > H2O
HO– > H2O > H3O+
H3O+ > H2O > HO–
2.12 Molecular Models


Organic chemistry is 3-D space
Molecular shape is critical in determining the chemistry a
compound undergoes in the lab, and in living organisms
84
2.13 Noncovalent Interactions

-
-
Several types:
Dipole-dipole forces
Dispersion forces
Hydrogen bonds
85
Dipole-Dipole
• Occur between polar molecules as a result of electrostatic
interactions among dipoles
• Forces can be attractive of repulsive depending on orientation
of the molecules
86
Dispersion Forces
• Occur between all neighboring molecules and arise
because the electron distribution within molecules
that are constantly changing
87
Hydrogen Bond Forces
• Most important noncovalent interaction in biological molecules
• Forces are result of attractive interaction between a hydrogen
bonded to an electronegative O or N atom and an unshared
electron pair on another O or N atom
88
89
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

90
Summary (cont’d)






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
Noncovalent interactions have several types: dipoledipole, dispersion, and hydrogen bond forces
91
Learning Check:
OH
O
+
phenol
pKa = 10
OH
hydroxide
+
H2O
phenoxide
pKa = 16
Which statement about the acid-base equilibrium
shown above is incorrect?
1.
2.
3.
4.
5.
the equilibrium favors the products
water is a conjugate acid in this equilibrium
hydroxide is the strongest base present
phenol is 6 times more acidic than water
the negative charge in the phenoxide is resonance
stabilized
Solution:
OH
O
+
phenol
pKa = 10
OH
hydroxide
+
H2O
phenoxide
pKa = 16
Which statement about the acid-base equilibrium
shown above is incorrect?
1.
2.
3.
4.
5.
the equilibrium favors the products
water is a conjugate acid in this equilibrium
hydroxide is the strongest base present
phenol is 6 times more acidic than water
the negative charge in the phenoxide is resonance
stabilized
Learning Check:
What is the most likely product of the following Lewis acidO
base reaction?
+
BF3
product
O
O
1.
2.
3.
4.
5.
A
B
C
D
E
BF3
O
BF3
O
O
O
BF3
D
O
O
O
BF3
A
B
C
O BF3
O
E
Solution:
What is the most likely product of the following Lewis acidO
base reaction?
+
BF3
product
O
O
1.
2.
3.
4.
5.
A
B
C
D
E
BF3
O
BF3
O
O
O
BF3
D
O
O
O
BF3
A
B
C
O BF3
O
E
Learning Check:
Which of these structures is not a resonance form of the
others? (The lone pairs are not shown explicitly.)
N
H
N
H
N
H
N
H
N
H
a)
b)
c)
d)
e)
1.
2.
3.
4.
5.
a
b
c
d
e
Solution:
Which of these structures is not a resonance form of the
others? (The lone pairs are not shown explicitly.)
N
H
N
H
N
H
N
H
N
H
a)
b)
c)
d)
e)
1.
2.
3.
4.
5.
a
b
c
d
e