Chemistry 343—Summer 2006
• General Information (Grading, Policies, etc.)
• Syllabus (Lectures, Quizzes, Exams)
• Recommended Problems
• Study Tips
• Chapter One: Basically Review (I hope);
Let’s Have at it…
Organic Chemistry: What and Why
• Compounds Based on Carbon
• Biological Molecules
• DNA
• RNA
• Amino Acids/Proteins
• Photosynthesis
• Pharmaceuticals
• A #&*$ Load of Other Stuff
Empirical vs. Molecular Formulas
• Empirical Formula: Lowest whole number ratio of
atoms in a given compound
• Molecular Formula: Exact composition of a
compound
Drawback: No Structural Information Provided by
Either
Later on we will look at methods that provide
structural detail
Empirical & Molecular Formula Examples
Consider 4 Hydrocarbons:
Ethene, Cyclopentane, Cyclohexane, 2-Butene
H
H
H
C
H
CH3
C
H3 C
H
H
Empirical Formula: CH2
Molecular Formula: C2H4, C5H10, C6H12, C4H8
Valence
• Valence best described as # of bonds an atom can form
Atom
Valence
Example
C
Tetravalent
CH4, CBr4
B, N
Trivalent
BH3, NH3
O
Divalent
H2O, H3C-O-CH3
H, Cl, Br
Monovalent
HCl, HBr, H2CCl2
• Related to # of valence electrons (Periodic Table)
Valence and the Periodic Table
• Valence Corresponds To Column (Group I, II, Nonmetals)
Increasing Electronegativity
Electronegativity and the Periodic Table
• Know the electronegativity trends!!
Lewis Structures
• Use only valence (outer shell) electrons
• Each atom acquires Noble gas configuration
• Octet Rule exceptions: Ions, Radicals, 3rd row
and lower (S, P, etc.)
• Sum # of valence electrons in atoms: this is the
number of electrons that should be represented
in the Lewis structure
• ½S(valence electrons) = # shared + lone pairs
Example: CH3Br
4 + 3(1) + 7 = 14 valence electrons
14/2 = 7 Shared/Lone pairs
H
H
C
H
Br
Example: C2H4
2(4) + 4(1) = 12 valence electrons
12/2 = 6 Shared/Lone Pairs
H
H
C
H
C
H
Example: CO324 + 3(6) + 2 = 24 valence electrons
24/2 = 12 Shared/Lone pairs
2O
C
O
O
• Place brackets around ions, indicate their charge
• We could have just as easily placed the double bond at other 2 O’s
Resonance: The Carbonate Ion
2-
2-
O
C
O
O
O
2-
O
O
C
C
O
O
O
• Double headed arrows indicate resonance forms
• Red “Curved Arrows” show electron movement
• Curved Arrow notation used to show electron flow in resonance
structures as well as in chemical reactions: we will use
this electron bookkeeping notation throughout the course
Octet Rule Exceptions: SO42• For now we focus on 3rd row atoms and beyond w/ ‘d’ orbitals
• Consider the sulfate ion: Here’s one valid Lewis structure
2O
6 + 4(6) + 2 = 32 valence electrons
O
S
O
32/2 = 16 Shared/Lone Pairs
O
• THIS IS NOT THE BEST POSSIBLE LEWIS STRUCTURE!
Formal Charge
• Formal Charge = #Valence Electrons - #Assigned Electrons
• We assign all electrons in a lone pair to an atom;
½ bonded electrons
2O
O
S
O
O
Formal Charges
S:
6 – 4 = +2
O:
6 – 7 = -1
• Lewis structures that minimize formal charge tend to be better
• Note: Sum of formal charges = molecular or ionic charge
d Orbitals & Minimizing Formal Charge
2O
6 + 4(6) + 2 = 32 valence electrons
O
S
O
32/2 = 16 Shared/Lone Pairs
O
_____Formal Charges_____
• Better Lewis structure with
minimized Formal Charge
• Note: There are resonance
structures (draw these?)
S
6–6=0
O(single)
6 – 7 = -1
O(double)
6–6=0
More Formal Charge Examples
_____Formal Charges_____
O
H
C
N
H
C:
O:
N:
H:
4–4=0
6–6=0
5–5=0
1–1=0
H:
N:
1–1=0
5–4=1
H
1+
H
H
N
H
H
Rules for Drawing Resonance Structures
2-
2-
O
C
O
O
O
2-
O
O
C
C
O
O
O
1. Hypothetical Structures; “Sum” Makes Real Hybrid Structure
2. Must be Proper Lewis Structures
3. Can Only Generate by Moving Electrons (NO Moving Atoms)
4. Resonance Forms are Stabilizing
5. Equivalent Resonance Structures Contribute Equally to Hybrid
Rules for Drawing Resonance Structures
6. More Stable Resonance Forms Contribute More to Hybrid
Factors Affecting Stability
1. Covalent Bonds
2. Atoms with Noble Gas (Octet) Configurations
H2C
O
CH3
vs.
H2C
O
CH3
3. Charge Separation Reduces Stability
4. Negative Charge on More Electronegative Atoms
Isomerism: Structural
• Structural Isomers: Same Molecular Formula; Different
Connectivity
• Why Might This Be a Big Deal? Consider Properties:
C2H6O
CH3CH2OH
CH3OCH3
BP
78.5 oC
-24.9 oC
MP
-117.3 oC
-138 oC
•Properties Can Differ Substantially Between Isomers!!
Isomerism: Cis/Trans
Cl
Cl
H
H
Cl
C
C
C
C
H
Cis or (Z)
Cl
H
Trans or (E)
• Same Molecular Formula (C2Cl2H2)
• Same Connectivity
• Different Structures  Double Bonds Don’t Rotate
Hybridization
For now, worry only about Carbon hybridization
Recall C’s valence configuration: 2s2 2p2
s orbital
p orbital
Will combine to form hybrid orbitals based on the valence
of the carbon atom
Hybridization (2)
Carbon Type
Hybridization
Alkane
sp3
Alkene
sp2
(one pure p left)
sp
(two pure p left)
Alkyne
Hybrid
Composition
25% s
75% p
33% s
67% p
50% s
50%p
Geometry
Tetrahedral
Trigonal
planar
Linear
Hybrid orbitals form single (s) bonds; pure p form multiple (p)
VSEPR Theory: What to Know
You are responsible for these geometries (the most
prevalent in Organic Chemistry):
Linear (e.g. acetylene)
Trigonal Planar (e. g. BF3, carbocations)
Trigonal Pyramidal (e.g. NH3, carbanions)
Tetrahedral (e.g. CH4, Ammonium Ion)
Angular (Bent) (e.g. H2O)
Representations of Organic Structures
• Condensed Formula: CH3CH2OH, CH3CH2CH2CH3
• Dash Formula:
H
H
H
C
C
H
H
O
• Bond-Line Formula
OH
H
H
H
H
H
H
C
C
C
C
H
H
H
H
H
Some Common Cyclic Structures
H 2C
H2C
CH 2
H2C
CH 2
H2
C
H2C
CH 2
H 2C
Cyclopropane
CH 2
H2C
Cyclobutane
H2
C
CH 2
Cyclopentane
H
C
H2C
CH 2
HC
CH
H2C
CH 2
HC
CH
C
H2
Cyclohexane
C
H
Benzene