Organic Chemistry I CHM 201

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Organic Chemistry I
CHM 201
William A. Price, Ph.D.
Introduction and Review:
Structure and Bonding
Atomic structure
Lewis Structures
Resonance
Structural Formulas
Acids and Bases
Electronic Structure of the Atom
• An atom has a dense,
positively charged nucleus
surrounded by a cloud of
electrons.
• The electron density is
highest at the nucleus and
drops off exponentially with
increasing distance from the
nucleus in any direction.
Chapter 1
5
Orbitals are Probabilities
2s Orbital Has a Node
The p Orbital
The 2p Orbitals
• There are three 2p
orbitals, oriented at
right angles to each
other.
• Each p orbital consists
of two lobes.
• Each is labeled
according to its
orientation along the x,
y, or z axis.
Chapter 1
9
px, py, pz
Electronic Configurations
• The aufbau principle states to
fill the lowest energy orbitals
first.
• Hund’s rule states that when
there are two or more
orbitals of the same energy
(degenerate), electrons will
go into different orbitals
rather than pairing up in the
same orbital.
Chapter 1
11
Electronic Configurations of Atoms
• Valence electrons are electrons on the
outermost shell of the atom.
Chapter 1
12
Covalent Bonding
• Electrons are shared between the atoms to complete
the octet.
• When the electrons are shared evenly, the bond is
said to be nonpolar covalent, or pure covalent.
• When electrons are not shared evenly between the
atoms, the resulting bond will be polar covalent.
Chapter 1
13
Lewis Dot Structure of Methane
carbon - 4 valence e
hydrogen - 1 valence e
.
.C.
.
H.
1s22s22p2
1s
Tetrahderal Geometry
Lewis Structures
CH4
NH3
H
Carbon: 4 e
4 [email protected] e ea: 4 e
8e
H C H
Nitrogen: 5 e
3 [email protected] e ea: 3 e
8e
H N H
H
H
H2O
Oxygen: 6 e
2 [email protected] e ea: 2 e
8e
Cl2
H O H
2 Cl @7 e ea: 14 e
Chapter 1
Cl Cl
16
Bonding Patterns
C
Valence
electrons
(group #)
4
# Bonds
# Lone Pair
Electrons
4
0
N
5
3
1
O
6
2
2
Halides
7
1
3
(F, Cl, Br, I)
Chapter 1
17
Bonding Characteristics of Period 2
Elements
Lewis structures are the way we
write organic chemistry.
Learning now to draw them
quickly and correctly will help
you throughout this course.
Chapter 1
20
Multiple Bonding
• Sharing two pairs of electrons is called a double bond.
• Sharing three pairs of electrons is called a triple bond.
Chapter 1
21
Convert Formula into Lewis Structure
•
•
•
•
•
•
•
•
HCN
HNO2
CHOCl
C2H3Cl
N2H2
O3
HCO3C3H4
Formal Charges
Formal charge = [group number ] – [nonbonding electrons ] – ½ [shared electrons]
H3O+
NO+
6 – 2 – ½ (6) = +1
6 – 2 – ½ (6) = +1
+
N O
H O H
H
+
5 – 2 – ½ (6) = 0
• Formal charges are a way of keeping track of electrons.
• They may or may not correspond to actual charges in the
molecule.
Chapter 1
23
Common Bonding Patterns
Chapter 1
24
Work enough problems to
become familiar with these
bonding patterns so you can
recognize other patterns as
being either unusual or wrong.
Chapter 1
25
Electronegativity Trends
Ability to Attract the Electrons in a Covalent Bond
Dipole Moment
• Dipole moment is defined to be the amount of charge
separation (d) multiplied by the bond length (m).
• Charge separation is shown by an electrostatic potential map
(EPM), where red indicates a partially negative region and
blue indicates a partially positive region.
Chapter 1
27
Methanol
Dipole Moment (m) is sum of the
Bond Moments
Nonpolar Compounds
Bond Moments Cancel Out
Nitromethane
Nitromethane has 2 Formal Charges
Form al C h arge = [Grou p #] - [# bon ds] - [# n on -bon di n g e
H
CH3NO2
H
H
O
C
N = 5-4-0 =+1
O = 6-1-6 =-1
N
O
Both Resonance Structures Contribute
to the Actual Structure
C H3NO2
H
H
H
C
O
H
O
H C
H
N
O
N
O
2 Equ i val e n t Re son an ce S tru ctu re s
Dipole Moment reflects Both
Resonance Structures
H
H
H
H
O
C
H C
N
H
O
O d
H
H
H
C
N
Od
Re son ance Hybrid
O
N
O
Resonance Rules
• Cannot break single (sigma) bonds
• Only electrons move, not atoms
3 possibilities:
– Lone pair of e- to adjacent bond position
• Forms p bond
 p bond to adjacent atom
 p bond to adjacent bond position
Curved Arrow Formalism
Shows flow of electrons
H
H
H
C
O
H
O
H C
H
N
O
N
O
Arrows de pict e le ctron pairs moving
Resonance Forms
• The structures of some compounds are not
adequately represented by a single Lewis structure.
• Resonance forms are Lewis structures that can be
interconverted by moving electrons only.
• The true structure will be a hybrid between the
contributing resonance forms.
Chapter 1
37
Resonance Forms
Resonance forms can be compared using the
following criteria, beginning with the most
important:
1. Has as many octets as possible.
2. Has as many bonds as possible.
3. Has the negative charge on the most
electronegative atom.
4. Has as little charge separation as possible.
Chapter 1
38
Two Nonequivalent Resonance
Structures in Formaldehyde
Major and Minor Contributors
• When both resonance forms obey the octet rule, the
major contributor is the one with the negative charge
on the most electronegative atom.
N C O
N C O
MAJOR
MINOR
The oxygen is more electronegative,
so it should have more of the negative
charge.
Chapter 1
40
Resonance Stabilization of Ions
Pos. charge is “delocalized”
H
H
C
H
C
C
H
H
C
H
H
H
C
C
H
H d
H
C C d
H
C
H
H
resonance hybrid
H
Solved Problem 2
Draw the important resonance forms for [CH3OCH2]+. Indicate which structure is
the major and minor contributor or whether they would have the same energy.
Solution
The first (minor) structure has a carbon atom with only six electrons around it. The
second (major) structure has octets on all atoms and an additional bond.
Chapter 1
42
Solved Problem 3
Draw the resonance structures of the compound below. Indicate which structure is
the major and minor contributor or whether they would have the same energy.
Solution
Both of these structures have octets on oxygen and both carbon atoms, and they
have the same number of bonds. The first structure has the negative charge on
carbon, the second on oxygen. Oxygen is the more electronegative element, so the
second structure is the major contributor.
Chapter 1
43
Resonance Forms for the
Acetate Ion
• When acetic acid loses a proton, the resulting acetate ion has
a negative charge delocalized over both oxygen atoms.
• Each oxygen atom bears half of the negative charge, and this
delocalization stabilizes the ion.
• Each of the carbon–oxygen bonds is halfway between a single
bond and a double bond and is said to have a bond order of
1½.
Chapter 1
44
Condensed Structural Formulas
Lewis
Condensed
H H
1
2
CH3CH3
H C C H
H H
• Condensed forms are written without showing all the
individual bonds.
• Atoms bonded to the central atom are listed after the central
atom (CH3CH3, not H3CCH3).
• If there are two or more identical groups, parentheses and a
subscript may be used to represent them.
Chapter 1
45
Drawing Structures
H
H
H
C
H
H
H
C
H
C
H
H
C
H
Bu tan e , C
4H10
CH3CH2CH2CH3
H
H
H
H C C C
H
H
H
C
H
H
H
Me th yl propan e , 4C
H10
CH
3CH(CH3)CH3
Octane Representations
C8H18 is molecular formula but there are 18 possible structural isomers
H
H H H H H H
H
C
C
C
C
H
C
C
C
C
H
H H H HH H
HH
Lewis structure
CH3CH2CH2CH2CH2CH2CH2CH3
CH3(CH 2)6CH3
condensed structural formula
Line-Angle Structures are Often
Used as a Short-hand
H
H H H H H H
H
C
C
C
C
H
C
C
C
C
H
H H H HH H
HH
Lewis structure
Line-angle or "zig-zag" structures
Line-Angle Structures
C12H26
H
H
H
H
H
C
H
H
H
C
H H
C
C
H
H
H
H
C
C
H
H
H
H
H
C
C
H
H
H
H
H
C
H
C
H
C
H
H
H
C
H
H
H
H
H
H H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
At the e nd of e ve ry l in e an d at th e in te rse ction of any l
th e re is acarbon atom with 4 bon ds. Hydroge n atom s
are m e n tall y su pplie d to fi ll th e vale ncy to 4.
H
H
H
H
H
H
H
H
H H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
Line-Angle structure Superimposed on
Lewis Structure
H
H
H
C
H
H
H
C
H H
C
C
H
H
H
H
C
H
C
H
H
C
C
H
H
H
H
C
H
C
H
C
H
H
C
H
H
H
Line-Angle Drawings
H H H H H
1
2
3
4
5
O
1
6
2
3
4
5
6
O
H C C C C C C
H H H H H
H
H
• Atoms other than carbon must be shown.
• Double and triple bonds must also be shown.
Chapter 1
52
For Cyclic Structures, Draw the
Corresponding Polygon
Cyclohexane
H H
H
H
C
C H
H C
H C
C H
C
H
H
H
H
Some Steroids
OH
O
HO
C h ol e ste rol
C27H44O
Te stoste ron e
C19H26O2
Definitions of Acids/Bases
+
Arrhe nius acid - form s3O
H in H2O
Arrhe nius base - forms Oin
H H2O
+
Bronste d-Lowry acid - donate s a (proton
H
)
+
Bronste d-Lowry base - acce pts a (proton
H
)
Le wis acid - acce pts an e le ctron pair to form a ne w bon d
Le wis base - don ate s an e le ctron pair to form a n e w bond
Dissociation in H2O
Arrhenius Acid forms H3O+
Bronsted-Lowry Acid donates a H+
Brønsted-Lowry Acids and Bases
Brønsted-Lowry acids are any species that donate a proton.
Brønsted-Lowry bases are any species that can accept a
proton.
Chapter 1
57
Conjugate Acids and Bases
• Conjugate acid: when a base accepts a proton, it becomes an
acid capable of returning that proton.
• Conjugate base: when an acid donates its proton, it becomes
capable of accepting that proton back.
Chapter 1
58
Acid Strength defined by pKa
HCl + H2O
H3O + Cl
[H3O ][Cl ]
Keq =
[HCl][H2O]
Ka = Keq[H2O] =
[H3O ][Cl ]
[HCl]
pKa = -log(Ka) = -7
7
= 10
Stronger Acid Controls Equilibrium
HCl + H2O
aci d
base
pKa = -7
stronger
H3O + Cl
con ju gate con ju gate
aci d
base
-1.7
weaker
Reaction Described with Arrows
H
H
Cl
+
O
H
H
+
O
H
H
Cl
Equilibrium Reactions
Identify the Acid and Base
O
CH3OH
+
OCOH
O
CH3O + HOCOH
Equilibrium Favors Reactants
O
O
CH3OH
aci d
pKa
15.5
+
OCOH
base
CH3O +
HOCOH
con j. base
con j. aci
6.5
The Effect of Resonance on pKa
Effect of Electronegativity on pKa
• As the bond to H becomes more polarized, H becomes more
positive and the bond is easier to break.
Chapter 1
66
Effect of Size on pKa
• As size increases, the H is more loosely held
and the bond is easier to break.
• A larger size also stabilizes the anion.
Chapter 1
67
Lewis Acids and Lewis Bases
• Lewis bases are species with available
electrons than can be donated to form a new
bond.
• Lewis acids are species that can accept these
electrons to form new bonds.
• Since a Lewis acid accepts a pair of electrons,
it is called an electrophile.
Chapter 1
68
Nucleophiles and Electrophiles
• Nucleophile: Donates electrons to a nucleus
with an empty orbital (same as Lewis Base)
• Electrophile: Accepts a pair of electrons (same
as Lewis Acid)
• When forming a bond, the nucleophile attacks
the electrophile, so the arrow goes from
negative to positive.
• When breaking a bond, the more
electronegative atom receives the electrons.
Chapter 1
69
Nucleophiles and Electrophiles
Chapter 1
70

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