CHEM 494
Special Topics in Chemistry
University of
Illinois at Chicago
UIC
Welcome to CHEM 494
Instructor: Prof. Duncan Wardrop
Time/Day: T, 4:30-8:15 p.m.
September 9, 2010
Course Website
http://www.chem.uic.edu/chem494
Syllabus
Course Policies
Other handouts
Announcements (Course News)
Course Calendar
University of
Illinois at Chicago
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CHEM 494, Fall 2012
Slide 2
Lecture 1
Keys to Success
1. Attend all lectures and discussion sections.
2. Don’t fall behind. Organic chemistry is easy, but each topic
builds on the previous.
•
3. Don’t memorize. Organic chemistry is conceptual.
•
4. Always ask yourself “Why” for everything you read or hear. You
may not always find the answer, but just asking will help you to find
connections and remember more.
University of
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UIC
CHEM 494, Fall 2012
Slide 3
Lecture 1
CHEM 494
Special Topics in Chemistry
University of
Illinois at Chicago
UIC
Origins and Examples of
Organic Chemistry
Introduction
Vitalism
Living
Nonliving.
vs.
•posses vital force
•compounds derived from are
“organic” (coined by J.
Berzelius, 1807)
•could not be synthesized in the
laboratory
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•termed “inorganic”
•derived from nonliving matter
•can be synthesized in the lab
CHEM 494, Fall 2012
Slide 5
Lecture 1
Wohler Synthesis Debunks Vitalism
F. Wohler
(Germany, 1828)
AgOCN(aq) + NH4Cl(aq)
X
Expected Product
H
H N H
H
O C N
ammonium cyanate
(inorganic)
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AgCl(s) + NH4OCN(aq)
Actual Product
H2N
O
C
NH2
urea
(organic)
CHEM 494, Fall 2012
Slide 6
Lecture 1
Self Test Question
What is the minimum requirement, today,
for a chemical substance to be classified
as organic?
A. derived from living matter
B. contains carbon
C. cannot be synthesized
D. bought at Whole Foods
E. combustion yields SO2
University of
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CHEM 494, Fall 2012
Slide 7
Lecture 1
Vitalism Lives On. . .
CH3
OH
HO
OH
OH
terpinen-4-ol
sold as tea tree oil
or Melaleuca oil
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O
OH
fructose
“fruit sugar”
also in HFCS
CHEM 494, Fall 2012
Slide 8
Lecture 1
OH
Organic Chemistry Everywhere
O
H3C
O
O
O
O
CH3
OH
O
O
OH
methyl salicylate
(oil of wintergreen)
3-methylbutyl acetate
(bananas)
H3C
N
O
UIC
O
O
cocaine
(analgesic)
University of
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aspirin
(analgesic)
N
O
procaine, "Novocaine"
(analgesic)
CHEM 494, Fall 2012
CH3
Slide 9
Lecture 1
Natural Products
O
O
N
H
C5H11
(-)-adalinine
A
6
3
NCS
HO
H2N
N
HN
HO
C
N
7
B
3'
C6H13
O
O
O
fasicularin
OH
OH
(-)-tetrodotoxin
University of
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CHEM 494, Fall 2012
Slide10
Lecture 1
Pharmaceuticals
CH3
N
F3C
Tamiflu
(influenza)
University of
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O
S NH2
O
Celebrex
(arthritis)
CHEM 494, Fall 2012
Morphine
(analgesic)
Slide 11
Lecture 1
Optics - Transition Lenses?
N
does not absorb
visible light
C
N
O
UV
light
heat
does absorb
visible light
N
C
N
University of
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CHEM 494, Fall 2012
O
Slide12
Lecture 1
CHEM 494
Special Topics in Chemistry
University of
Illinois at Chicago
UIC
Atomic Structure
A General Chemistry
Review
Atomic Composition
particle
-
++
+
-
+
proton
neutron
-
electron
charge
mass
molar mass
symbol
positive
1.6726 x 10-24 g
1.0073 g/mol
p
neutral
1.6750 x 10-24 g
1.0087 g/mol
n
negative
9.1096 x 10-38 g
5.486 x 10-4 g/mol
e–
Bohr Model
University of
Illinois at Chicago
UIC
CHEM 494, Fall 2012
Slide14
Lecture 1
Atomic Number & Mass Number
A
Z
X
A = protons + neutrons
Z = protons
6
Li
-
++
+
-
University of
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3
For neutral molecules, the number of
electrons equals the number of
protons.
CHEM 494, Fall 2012
Slide15
Lecture 1
Self Test Question
How many neutrons are
in the following atom:
14
6
C
A. 14
B. 6
C. 8
D. 20
E. cannot be determined
University of
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UIC
CHEM 494, Fall 2012
Slide16
Lecture 1
Tenets of Schrӧdinger Equation
•
•
•
electrons have wave properties
wave equation gives energy of electron at a location
solutions to wave equation are wave functions (Ψ); a.k.a
orbitals
• probability distribution = Ψ2 (Heisenburg uncertainty
principle)
University of
Slide17
Illinois at Chicago
UIC
CHEM 494, Fall 2012
Lecture 1
Probability Distribution vs. Boundary
Surface
1s
1s
Probability distribution (Ψ2)
(“electron cloud”)
Boundary Surface (1s orbital)
(where the probability = 9095%)
University of
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CHEM 494, Fall 2012
Slide18
Lecture 1
Wave Function Values:
Quantum Numbers
• Shrodinger equation → wave function (orbital, Ψ)
• many solutions for Ψ, energies of electrons in atom
quantum
number
principle
orbital
magnetic
spin
symbol
n
l
ml
ms
values
examples/
abbreviations
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1, 2, 3. . .
(ml = −l, −l+1 ... 0
0, 1, . . . n-1
+1/2 or -1/2
... l−1, l)
l= 0, s
l = 1, p
l = 2, d
l = 3, f
CHEM 494, Fall 2012
s: 1 orbital
p: three orbitals
d: five orbitals
f: seven orbitals
Slide19
Lecture 1
s-Orbitals
•
spherically symmetric
• possible for n≥1
•
1 s-orbitals for each value of
n
• 1s = no nodes, 2s = 1
nodes, 3s = 2 nodes, etc.
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•
probability of finding selectron at nodal surface =
0
•
s-orbital energy increases
with increasing nodes (as n
increases)
CHEM 494, Fall 2012
Slide20
Lecture 1
p-Orbitals
•
shaped like dumbells
• not possible for n = 1 (n≥2)
•
3 p-orbitals for each value
of n; they are degenerate
(equal in energy)
•
wave function changes sign
at the nucleus (node)
• probability of finding pelectron at nodal plane = 0
•
University of
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CHEM 494, Fall 2012
higher in NRG than sorbitals of the same shell
Slide21
Lecture 1
Relative Energies of Orbitals
n = 3/4
3d 3d 3d 3d 3d
4s
3p 3p 3p
n=2
2p 2p 2p
2s
n=1
University of
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Energy
3s
1s
UIC
CHEM 494, Fall 2012
Slide22
Lecture 1
Order of Orbitals from Periodic Table
1s
2s
3s
4s
5s
6s
7s
Read left to right starting at the top
left; just like a typewriter.
2p
3p
4p
5p
6p
7p
3d
4d
5d
6d
4f
5f
University of
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CHEM 494, Fall 2012
Slide23
Lecture 1
Electron Configuration: Filling Orbitals
3d 3d 3d 3d 3d
4s
3p 3p 3p
3s
spin (ms = +1/2)
n=2
2p 2p 2p
2s
n=1
University of
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1s
UIC
spin (ms = –1/2)
Energy
n = 3/4
1. Aufbau principle: fill lowest energy orbitals first.
2. Pauli exclusion principle: two electrons with same
quantum numbers cannot occupy a single orbital.
Electrons must have opposite spins in the same orbital.
3. Hund’s rule: degenerate orbitals filled singly first.
CHEM 494, Fall 2012
Slide24
Lecture 1
Self Test Question
Which of the following orbitals is
highest in energy?
A. 3f
B. 3p
C. 2s
D. 3s
E. 2d
University of
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CHEM 494, Fall 2012
Slide25
Lecture 1
Self Test Question
Which is the correct electron
configuration for carbon?
A. 1s2, 2s8
B. 1s2, 2s4
C. 1s2, 2s2, 2s2
D. 1s2, 2s2, 2p2
E. none of the above
University of
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UIC
CHEM 494, Fall 2012
Slide26
Lecture 1
CHEM 494
Special Topics in Chemistry
University of
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Bonding and Molecular
Structure
What is a Covalent Chemical Bond?
valent = bonding
covalent = electrons shared between two nuclei
Forces involved:
•electron-electron
repulsion
•nucleus-nucleus
repulsion
•electron-nucleus
attraction
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electron-electron repulsion < electron-neutron attraction
CHEM 494, Fall 2012
Slide28
Lecture 1
Three Models of Bonding
Lewis Model:
•atoms gain, lose or share electrons in order to achieve a closed-shell
electron configurations (all orbitals fully occupied)
•closed-shell for row 2 = 8 electrons (octet rule)
•only valence electrons (in outermost shell) are involved in bonding
Atomic Structures
Molecular
Lewis Structures
Molecular
Lewis Structures
H–H
F–F
University of
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CHEM 494, Fall 2012
Slide29
Lecture 1
Three Models of Bonding
Valence Bond Model:
•in-phase overlap of two half-filled orbitals
•in-phase = constructive interference
•increases probability electrons between two
nuclei
Boundary
Surfaces
H
H
H–H
Electrostatic
Potential maps:
red = negative
blue = postive
University of
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CHEM 494, Fall 2012
Slide30
Lecture 1
Three Models of Bonding
Molecular Orbital (MO) Model:
•combine atomic orbitals (AO) of all atoms, then extract molecular
orbitals
•number of atomic orbitals in equals the number of molecular orbitals out
•combination possibilites:
•additive = produces bonding molecular orbitals
•subtractive = produces antibonding molecular orbitals
University of
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CHEM 494, Fall 2012
Slide31
Lecture 1
Molecular Orbital Diagram
nodal plane
e- probability = 0
= no bond
antibonding (σ*)
MO
H
H
1s AO
1s AO
Rules Still Apply:
1. Aufbau
2. Pauli
3. Hund
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bonding (σ) MO
CHEM 494, Fall 2012
Slide32
Lecture 1
Self Test Question
Which is the correct Lewis dot structure
for CH4O?
A.
H
H C H O H
H
B.
H
C.
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O
C
H
D.
H
H C O H
H
O H
H
H
H
H C
E. none of
the above
CHEM 494, Fall 2012
Slide33
Lecture 1
Structural, Bond-line and Condensed
Formulas
Structural
•format varies; typically most
covalent bonds are drawn out
•bonds are lines
•only lone-pair (nonbonding)
electrons are drawn as dots
H H O
H
C
C
C
H
C
H HH H
H O H
H
Bond-line
Condensed
•only atoms written are those
•all or most covalent bond lines
that are not C or H bound to C
•intersection of two lines is C
•terminus of a line is -CH3
group
•# H atoms is assumed for C
•chains drawn as “zig-zag”
are omitted
•groups are separated by
parentheses (infers group is
attached to carbon with
available valency on left or
right)
CH3CH2CH2CHO
OH
H
C
H HH H
H
C
C
CH3CH(OH)CH3
HH
HH H
H
C
C
O
C
C
C
H
H HH H O
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O
H
(CH3)2CHCH2CH2CO2H
O
CHEM 494, Fall 2012
Slide34
Lecture 1
VSEPR: Quick Review
University of
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CHEM 494, Fall 2012
Slide35
Lecture 1
Self Test Question
Which bond-line drawing correctly
represents the following condensed
formula?
CH3CH(OCH2CH3)CH2Br
O
A.
Br
D.
O
Br
O
Br
B.
E.
O
Br
O
C.
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Br
CHEM 494, Fall 2012
Slide36
Lecture 1
CHEM 494
Special Topics in Chemistry
University of
Illinois at Chicago
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Functional Groups
Memorize These Functional Groups NOW
functional group: a defined
group of atoms with a specific
connectivity
•responsible for properties
•predictable reactivity
•well-defined nomenclature
Memorizing: “What is the minimum
number of atoms needs to define a
functional group and in what order
are they bonded?” Flashcards!
Table 4.1, page 140
University of
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CHEM 494, Fall 2012
Slide38
Lecture 1
Formal Charge - Example 1
My logic (no memorizing equations):
1.Recognize that opposing charges cancel each other
2.Determine difference in protons vs. electrons by asking:
a.How many valence (outer shell) electrons does the atom have?
b.How many valence (outer shell) electrons does the atom “want” (groups #)?
1.If atom has more than it “wants,” negatively charged.
2.If atom has less than it “wants,” positively charged.
2p
2s
-
1s
+
+ +
+
+
+ +
+
-
16
-
-
-
-
2–
O
8
Has? = 8
Wants? = 6
Difference? =2 extra
-
University of
Illinois at Chicago
UIC
CHEM 494, Fall 2010
Slide 39
Lecture 2
Formal Charge - Example 2
My logic (no memorizing equations):
1.Recognize that opposing charges cancel each other
2.Determine difference in protons vs. electrons by asking:
a.How many valence (outer shell) electrons does the atom have?
b.How many valence (outer shell) electrons does the atom “want” (groups #)?
If atom has more than it “wants,” negatively charged.
If atom has less than it “wants,” positively charged.
H
Has? = 5
O
H
H
“formal” = count one electron for each
bond
University of
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CHEM 494, Fall 2010
Wants? = 6
Difference? =1 less
Slide 40
Lecture 2
Shortcut for Some 1st and 2nd Period Atoms
Atom
Bonds
Formed
Group
H
1
I
C
4
IV
N
3
V
O
2
VI
F (X)
1
VII
University of
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• For each additional bond =
+1
• For each “missing” bond = –
1
CH3
C
H 3C
CH3
N
O
H3C
O
CHEM 494, Fall 2010
Slide 41
Lecture 2
Self Test Question
What is the formal charge on the
carbon atom in red?
H
H
H
C
H
C
H
A. -2
B. -1
C. 0
D. +1
E. +2
University of
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CHEM 494, Fall 2010
Slide 42
Lecture 2
Self Test Question
According to VSEPR, what is the
bond angle between 2 H-atoms in
methane (below)?
H
H
University of
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?
C
H
H
CHEM 494, Fall 2010
A. 105º
B. 109.5º
C. 107º
D. 180º
E. 120º
Slide 43
Lecture 2
CHEM 494
University of
Special Topics in Chemistry Illinois at Chicago
UIC
Structure and
Bonding: Resonance
Section: 1.8, 1.11
You are responsible for Section 1.10.
The Contradictory Case of Ozone
Lewis Structure
suggests ozone is
non-symmetrical
electrostatic potential map
red = negative
blue = positive
O
O
O
Microwave spectroscopy shows
that ozone is symmetrical (C2v)
University of
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CHEM 494, Fall 2010
Slide 45
Lecture 2
Curious Case of Benzene
Predicted
Actual
H
H
H
resonance
H
H
H
C–C bond length: 150 pm
C=C bond length: 134 pm
University of
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All bonds = 140 pm
CHEM 494, Fall 2010
Slide 47
Lecture 2
Resonance Provides the Solution
Problems:
1.Lewis structures fail to describe actual atom electron densities for
molecules with more than one possible electron distribution.
2.Lewis formulas show electrons as localized; they either belong to a
single atom (lone pair) or are shared between two atoms (covalent
bond).
3.Electrons are not always localized; delocalization over several nuclei
leads to stabilization (lower energy).
O
H3C
University of
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O
O
CH3
N
CH3
UIC
H3C
CH3
N
CH3
CHEM 494, Fall 2010
H3C
CH3
N
CH3
Slide 48
Lecture 2
Physical Proof of Resonance in Amides
O
H 3C
O
N
CH3
H3C
CH3
N
CH3
CH3
C-N (amide) Bond Length = 133 pm
C-N (amine) Bond Length = 147 pm
C=N (imine) Bond Length = 129 pm
H
N
N
H
H
H
University of
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H3C
UIC
CHEM 494, Fall 2010
CH3
Slide 49
Lecture 2
A Question of Arrows
Equilibrium between distinct species
Reaction from one species to another
“Movement” of electrons
from donor to acceptor
Indicates that 2 species are
contributing resonance
structures
University of
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CHEM 494, Fall 2010
Slide 50
Lecture 2
Resonance
Solutions:
1.Actual molecule is considered a resonance hybrid (weighted average)
of contributing Lewis structures. (Note: double headed arrow does NOT
indicate interconversion.)
2.Dashed-line notation is sometimes used to indicate partial bonds.
3.Not all structures contribute equally.
University of
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UIC
CHEM 494, Fall 2010
Slide 51
Lecture 2
Rules of Resonance
You are responsible for the “Rules of Resonance” (Pg. 27)
• How to draw resonance structures
• How to distinguish between a resonance
structure and a unique electron configuration
• How to predict the major contributing structure
University of
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UIC
CHEM 494, Fall 2010
Slide 52
Lecture 2
Another Way to Think
About Resonance Contributors
Dürer's Rhinoceros
(1515 A.D.)
University of
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UIC
CHEM 494, Fall 2010
Slide 53
Lecture 2
Curved Arrow Notation
curved arrows show
the movement of
electrons; never
atoms
O
CH3
N
CH3
H 3C
O
H3C
CH3
N
CH3
University of
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atoms
electrons
resonance: electrons
in a covalent bond
moving out to an atom
resonance: lone
pair of electrons
moving in between
two atoms to form a
new covalent bond
UIC
H
O
bond making: lone
pair of electrons
+
H3O forming a new bond
to another atom
H
H
O
H3C
O
O
H
CHEM 494, Fall 2010
H3C
O
+ H
bond breaking:
electrons in a bond
leaving to most
electronegative atom
Slide 54
Lecture 2
Preview: Curved Arrows in Reaction
Mechanisms
bond making
O
H3C
O
O
H
+
H
O
H
H3C
O
+
H
H
O
bond breaking
University of
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UIC
CHEM 494, Fall 2010
Slide 55
Lecture 2
H
Another Way of Viewing Arrow-Pushing
Remember that curly arrows originate from
e-donors (nucleophiles) and terminate at eacceptors (electrophiles)
From an M.O. perspective curly arrows
originate from filled orbitals (nucleophiles)
and
terminate
at
empty
orbitals
(electrophiles) - bonds are formed from the
mixing of filled and empty orbitals
University of
Illinois at Chicago
UIC
CHEM 494, Fall 2010
Slide 56
Lecture 2
Molecular Orbital Diagram
nodal plane
e- probability = 0
= no bond
antibonding (σ*)
MO
H
H
1s AO
1s AO
Rules Still Apply:
1. Aufbau
2. Pauli
3. Hund
University of
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bonding (σ) MO
CHEM 494, Fall 2010
Slide 57
Lecture 2
Self Test Question
Which of the following is not
descriptive of resonance as it
pertains to organic chemistry?
A.
B.
C.
D.
E.
University of
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stabilization of negative charge
movement of atoms
delocalization of electrons
involves bond-breaking
involves bond-making
UIC
CHEM 494, Fall 2010
Slide 58
Lecture 2
CHEM 494
University of
Special Topics in Chemistry Illinois at Chicago
UIC
How Structure Affects
Physical Properties:
Acid Strength
Arrhenius Acids and Bases
Acid: Dissociates to provide
protons (H+) in water.
strong acid: ionizes
completely
weak acid: does not ionize
completely
Base: Dissociates to
provide hydroxide (OH–) in
water.
strong base: ionizes
completely
weak base: does not ionize
completely
University of
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CHEM 494, Fall 2010
Slide 60
Lecture 2
Brønsted-Lowry Acids & Bases
Acid: Proton (H+) donor
Base: Proton (H+) acceptor
B:
+
Base
H
Acid
A
B H
Conjugate Acid
A:
+
Conjugate Base
• definition does not depend on dissociation in water
• broader definition
• conjugate acid and base differ by only one proton
• most often used definition in this course
University of
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CHEM 494, Fall 2010
Slide 61
Lecture 2
Molecular Orbital Picture
H-A Bond Breaks!
B
:
H
A
B
H
n (B)
* (H-A)
 (B-H)
Non-Bonding M.O.
Anti-Bonding M.O.
Bonding M.O.
University of
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CHEM 494, Fall 2010
+
A
n (A)
Non-Bonding
M.O.
Slide 62
Lecture 2
:
Water Can Act as a Base or Acid Under
Brønsted-Lowry Definition
H
O
H
+
water
(base)
Cl
HCl
(acid)
H
N
H
H
O
H
amide
(base)
University of
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H
water
(acid)
UIC
H
H
O
H
+
Cl–
hydronium ion
chloride ion
(conjugate acid) (conjugate base)
H
H
N
H
+
OH–
ammonia
hydroxide ion
(conjugate acid) (conjugate base)
CHEM 494, Fall 2010
Slide 63
Lecture 2
Acid Strength is Measured by the Acid
Dissociation Constant (pKa)
H
Acid
A
H
Proton
+
A:
Anion
pKa = –log10 Ka
Note that the value of K reflects the relative
thermodyamic stability of acid and anion
University of
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UIC
CHEM 494, Fall 2010
Slide 64
Lecture 2
pKa Scale
acid
pKa
HCl
hydrochloric
-3.9
H3O+
hydronium
-1.7
acetic acid
4.7
O
H3C
OH
NH3
HA
pyridinium
5.2
OH
phenol
10
H2O
water
15.7
CH3CH2O H
ethanol
16
NH3
ammonia
36
CH3CH3
ethane
62
University of
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Take home message:
the larger the Ka value,
the smaller the pKa value,
the stronger the acid.
CHEM 494, Fall 2010
acid
dissociation
H
+
A–
conjugate
base
[H+][A-]
pKa = -log10 [HA]
-log10(1/100) = 2
-log10(1/10) = 1
-log10(1) = 0
-log10(10) = -1
-log10(100) = -2
Slide 65
Lecture 2
Structure Affects Acid Strength
1. Electronegativity of H-Donating Atom
•dominant effect for same periods (rows)
•more electronegative conjugate base = more stable conjugate base
= Ka lies further to right
•Alternative reasoning: H of conjugate acid (HA) becomes more
positive with increasing electronegativity of A
least stable
(highest
energy)
conjugate base
= weakest acid
most stable
(highest
energy)
conjugate base
= strongest acid
acid
conjugate base
University of
Illinois at Chicago
CH3
UIC
NH2
CHEM 494, Fall 2010
OH
F
Slide 66
Lecture 2
Structure Affects Acid Strength
1. Electronegativity (cont.)
•inductive effect of electronegative substituents also stabilizes conj.
base
•the more electronegative the group, the greater the stabilization
•the closer the electronegative group, the greater the inductive effect
•inductive effect = structural effects that are transmitted through
bonds
H H
H H
Ka = 1 x 10-16
+
H C C O H
H C C O
H
pKa = 16
H H
H H
F H
F C C O H
F H
University of
Illinois at Chicago
UIC
F H
F C C O
F H
+
H
CHEM 494, Fall 2010
Ka = 5 x 10-12
pKa = 11.3
Slide 67
Lecture 2
Self Test Question
Which of the following carboxylic
acids (-CO2H) is the most acidic
(lowest pKa)?
Cl
O
C
B.
H
O
C
C.
H H O
C
H
Cl
C
C
C
O
H HH H
A.
O
R
O
carboxylate anion
(conjugate base)
University of
Illinois at Chicago
UIC
C
H H
C
H H
O
O
H
H
CHEM 494, Fall 2010
D.
E.
Cl
O
C
F
O
C
C
H H
C
H H
O
O
H
H
Slide 68
Lecture 2
Structure Affects Acid Strength
2. Bond Strength
•dominant effect for same groups (columns)
•Similar sized bonding orbitals = better overlap = stronger
bond
•Stronger H–A bond = weaker acid
University of
Illinois at Chicago
UIC
CHEM 494, Fall 2010
Slide 69
Lecture 2
Self Test Question
Which of the following is the
strongest acid?
University of
Illinois at Chicago
UIC
A. H2O
O
B. H2S
S
CHEM 494, Fall 2010
Slide 70
Lecture 2
Structure Affects Acid Strength
3. Resonance (delocalization of e–s in conj. base):
•delocalization of e–s in conj. base = increased stability (lower
NRG) of conjugate base
•more stable conj. base = larger Ka = smaller pKa = stronger acid
H H
C
H
H3C
O
H3C
O
C
O
H
H H
C
H3C
O
H3C
H3C
University of
Illinois at Chicago
UIC
O
C
O
C
O
+
+
O
H
Ka = 1 x 10-16
pKa = 16
H
Ka = 2.0 x 10-5
pKa = 4.7
H3C
CHEM 494, Fall 2010
O
C
O
Slide 71
Lecture 2
Acid-Base Equilibria: Determining the
Direction of Acid-Base Reactions
You must identify the ACID on each side of the equilibrium:
pKeq = pKa (acid left) - pKa (acid right)
-[pKa (acid left) - pKa (acid right)]
Keq = 10
• remember: p = -log10
• this equation works for any acidbase reaction; doesn’t matter
which way equilibrium is written
Example:
O
H
+
phenol
(pKa = 10)
acid
HO
O
C
O
Keq
+
O
hydrogen carbonate
(pKa = 10.2)
base
phenoxide anion
(pKa = NA)
base
-[10 - 6.4]
-[3.6]
Keq = 10
= 10
HO
O
C
OH
carbonic acid
(pKa = 6.4)
acid
= 2.5 x 10-4
since Keq <1, then equilibrium lies to the left
University of
Illinois at Chicago
UIC
CHEM 494, Fall 2010
Slide 72
Lecture 2
Finally, What About Strengths of Bases?
Question: Is
ammonia or pyridine
a stronger base?
Solution:
1. Determine which conjugate acid
of each base is the weakest.
2. The weaker the conjugate acid,
the stronger the conjugate base.
University of
Illinois at Chicago
UIC
CHEM 494, Fall 2010
Slide 73
Lecture 2
Molecule of the Week...α-Amanitin
Be Careful What You Eat (and Who You
Marry)!
Amanita phalloides
α-Amanitin
α-Amanitin (red) bound
to RNA polymerase II
Tiberius Claudius Caesar Augustus Germanicus (10 B.C.–A.D. 54)
Read more about the death of Claudius......
University of
Illinois at Chicago
UIC
CHEM 494, Fall 2010
Amanita caesarea
Slide 74
Lecture 2