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Organic Chemistry CHM 314 Dr. Laurie S. Starkey, Cal Poly Pomona
Structure & Shape of Organic Molecules - Chapter 1 (Wade)
Organic (living things, chemistry of carbon)
Inorganic (rocks, minerals, metals, glass)
Examples of Organic Compounds:
Products
medicine
pesticides
dye/paint/ink
gasoline/fuels
cosmetics
Materials
paper
cotton
tires/rubber
nylon/polyester
plastic/vinyl
in Nature
hormones/steroids
DNA
protein/fats/sugars
flavors/fragrances
molecular bio. = organic rxns
Review Some Chemistry Basics (Wade 1-1 to 1-8)
electronic configurations and the Periodic Table
Energy
• electrons (e-) are held in atomic orbitals
around the nucleus (s, p, d, f), and s orbitals are
more stable (lower Energy) than p orbitals
2p ____ ____ ____
2s ____
1s ____
• fluorine is the most electronegative element
(pulls electron density toward itself)
Which is more electronegative, C or N?
Periodic
trends:
F
• oxygen is the second-most electronegative
element and C ≈ H (no significant difference)
• all elements want to "look like" the Noble
gases (have same electronic configuration)
atom is stable if it has a filled valence shell
(octet rule)
Ex: Na+ Ca2+ Br–
8 e– = s2p6 (or 1s2 for He)
ionic bonds are formed between atoms if they have a large difference in electronegativities
Na
Cl
transfer e-
unstable, reactive
Na
Cl
a salt
(network
structure)
stable, filled octets
covalent bonds are formed between atoms if they have similar electronegativities
H
H
C
H
H
unstable, reactive
share
e-
H
H C H
H
H
=
H C H
H
stable, filled octets
methane
molecule
represents
2 shared e(covalent bond)
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polar covalent bonds arise if there is a difference in electronegativities between atoms
C
C
O
H
nonpolar bond
(there is no significant difference
in electronegativities)
polar bond
(O is more elcgtronegative than C)
polarity of molecule depends on geometry (more in Wade Chapter 2)
O C O
linear molecule: equal and opposite polar bonds, so
no net dipole moment
nonpolar
CH3
O
CH3
bent molecule: has a net dipole moment
polar
CH3
CH2
no polar bonds: no net dipole possible
(regardless of geometry)
CH3
nonpolar
ionic: full charges are the ultimate in polarity
hydroxide (OH-) is called a covalent ion
NaOH has both ionic and covalent bonds
Na
O H
(NaOH)
ionic
!line drawings (Wade 1-10) are a short-hand way to draw carbon structures
• end points and intersections represent C atoms
• omit H’s attached to C’s
HC
CH3CH2CH2CH2CH2CH2CH2CH3
HC
Octane
(gasoline component)
H
C
CH
CH
C
H
Benzene
CH3
OH
CH3
CH OH
O
Isopropyl Alcohol
(rubbing alcohol)
O
CH2CHCHO
O
OH
CH3
C
OH
Acetic acid
(vinegar)
(CH3)3CCCCHCl2
Cl
CH3CO2H
Cl
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Drawing Lewis Structures (Wade 1-4 to 1-8)
Drawing Lewis Structures
1) draw skeleton - connectivity
example ClCH2CN
2) count total # of valence electrons
(valence e– = group no.)
3) subtract charge (if any)
4) fill in missing electrons (fill octets)
5) determine formal charges (if any)
example CH3OH2+
Formal Charges
• calculate for each atom
• determine "electron count"
= all nonbonded + 1/2 bonded/shared
• compare "electron count" with valence
missing an electron −−> + charge
extra electron −−> – charge
Typical, stable bonding (know by inspection)
Atom
H
C
N
O
X
X = halogen
(F, Cl, Br, I)
example
# bonds
# lone pairs
"e- count"
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practice: Draw the Lewis structure of nitric acid, HNO3
Understanding Resonance (Wade 1-9)
1. Any compound for which more than one Lewis structure may be written is accurately described
by no single structure. The actual structure is a resonance hybrid of them all (NOT "flipping back
and forth" between resonance forms). The various structures are called contributing structures or
resonance forms.
2. The stability of a resonance hybrid is greater than that which would be expected of any of the
contributing structures. The hybrid is therefore said to be stabilized by resonance or resonance
stabilized (by an amount of energy called the resonance energy).
3. The greater the stability (the lower the energy) of a contributing structure, the greater will be its
contribution to the total structure of the hybrid.
4. Resonance is the interaction of electrons in p orbitals. Only pi and nonbonded electron density is
reorganized in resonance (no σ bonds break). No atoms move (no change in bond length, angles).
Since electrons are not being added or removed, there is no change in overall charge.
Rules for Estimating Stability of Resonance Structures
1. The greater the number of filled octets, the greater the stability. You might notice that a
structure with more covalent bonds is more stable (since more atoms will have complete octets).
CH3
O
CH
CH3
CH3
less important
O
CH
CH3
more important
2. The structure with fewer formal charges is more stable. If the species already has a net charge,
creation of new charges is not favorable. For such charged compounds, the goal in drawing
resonance forms is to delocalize the charge – relocate it to as many different positions as possible.
O
H
C
O
NH2
H
more important
(best Lewis structure)
C
O
NH2
H
less important
C
NH2
least important
3. Other things being equal, a structure with a negative charge on the more electronegative
element will be more stable. Similarly, a positive charge on a less electronegative element is
more stable.
H
H
H
H
C
H
O
C
H
important resonance
O
C
H
O
C
H
not important resonance (unlikely)
O
important resonance
not important resonance (unlikely)
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4. Resonance forms that are equivalent have no difference in stability and contribute equally.
CH2
CH
CH2
CH2
CH
CH2
"allyl" carbocation
both forms are
equal in energy
Looking for Resonance Delocalization of Electrons (3 main types/patterns*)
1. lone pair next to a pi bond (allylic lone pair)
CH2
CH
CH2
2. vacancy (missing octet) next to a pi bond (allylic carbocation)
CH2
CH
CH2
3. pi bond between two different elements (carbonyl-like resonance)
**electrons move toward the more electronegative element**
O
H
C
H
*Note: we will encounter a fourth
important type of resonance later,
called aromatic resonance, that
is found in molecules like benzene
benzene
practice: draw and rank all possible resonance forms (explain rankings)
O
HO
C
NH2
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Acid-Base Reactions: Proton Transfers (Wade 1-13)
Topic Outline
1) Definitions (1-12 to 1-14)
2) Periodic Trends
3) Inductive Effects
4) Resonance Effects
5) Common Acids and Bases
See also (at the end of the books):
Wade Appendices 4 and 5 for pKa Tables
and Solutions Manual Appendix 2,
"Summary of Acidity and Basicity"
1) Definitions: acids and bases can be defined by Lewis or Bronsted-Lowry theories
Lewis Acid: electron-pair acceptor (also called an Electrophile, E+)
* has a vacancy
* common Lewis acids:
AlCl3
BF3
Lewis Base: electron-pair donor (also called an Nucleophile, Nu:)
* has a lone pair or a pi bond
examples:
CH3
C
NH3
AlCl3
base
"acid"
CH3
H
O
H
H
"Acid-Base" reaction usually means Bronsted-Lowry type
Acid: H (proton) donor
Base: H
(proton) acceptor
(Bronsted-Lowry
definitions)
A general "proton-transfer" reaction
H
B
base
A
acid
Two acids are in competition - forward and reverse reactions are in equilibrium.
**Equilibrium lies in the direction of the ______________ acid/base pair **
Which is the stronger acid? Use pKa table (Wade Appendix 4 and 5) or predict...
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