Transcript Organic Chemistry: What and Why
Chemistry 343—Summer 2006
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General Information (Grading, Policies, etc.)
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Syllabus (Lectures, Quizzes, Exams)
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Recommended Problems
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Study Tips
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Chapter One: Basically Review (I hope); Let’s Have at it…
Organic Chemistry: What and Why
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Compounds Based on Carbon
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Biological Molecules
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DNA
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RNA
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Amino Acids/Proteins
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Photosynthesis
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Pharmaceuticals
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A #&*$ Load of Other Stuff
Empirical vs. Molecular Formulas
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Empirical Formula: Lowest whole number ratio of atoms in a given compound
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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 C H H H 3 C CH 3 H Empirical Formula: CH 2 Molecular Formula: C 2 H 4 , C 5 H 10 , C 6 H 12 , C 4 H 8
Valence
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Valence best described as # of bonds an atom can form Atom
C B, N O H, Cl, Br
Valence
Tetravalent Trivalent Divalent Monovalent
Example C
H 4 ,
C
Br 4
B
H 3 ,
N
H 3 H 2
O
, H 3 C-
O
-CH 3 HCl, HBr,
H
2 C
Cl
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Related to # of valence electrons (Periodic Table)
Valence and the Periodic Table
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Valence Corresponds To Column (Group I, II, Nonmetals)
Electronegativity and the Periodic Table
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Know the electronegativity trends!!
Lewis Structures
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Use only valence (outer shell) electrons
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Each atom acquires Noble gas configuration
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Octet Rule exceptions: Ions, Radicals, 3 rd row and lower (S, P, etc.)
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Sum # of valence electrons in atoms: this is the number of electrons that should be represented in the Lewis structure
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½
S
(valence electrons) = # shared + lone pairs
Example: CH 3 Br 4 + 3(1) + 7 = 14 valence electrons 14/2 = 7 Shared/Lone pairs H H C H Br
Example: C 2 H 4 2(4) + 4(1) = 12 valence electrons 12/2 = 6 Shared/Lone Pairs H C H C H H
Example: CO 3 2 4 + 3(6) + 2 = 24 valence electrons 24/2 = 12 Shared/Lone pairs 2 O C O O
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Place brackets around ions, indicate their charge
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We could have just as easily placed the double bond at other 2 O’s
Resonance: The Carbonate Ion O O C O 2 O O C O 2 O O C O 2-
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Double headed arrows indicate resonance forms
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Red “ Curved Arrows ” show electron movement
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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: SO 4 2-
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For now we focus on 3 rd row atoms and beyond w/ ‘d’ orbitals
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Consider the sulfate ion: Here’s one valid Lewis structure
2 O O S O O 6 + 4(6) + 2 = 32 valence electrons 32/2 = 16 Shared/Lone Pairs
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THIS IS NOT THE BEST POSSIBLE LEWIS STRUCTURE!
Formal Charge
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Formal Charge = #Valence Electrons - #Assigned Electrons
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We assign all electrons in a lone pair to an atom; ½ bonded electrons 2 O Formal Charges S: 6 – 4 = +2 O S O O: 6 – 7 = -1 O
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Lewis structures that minimize formal charge tend to be better
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Note: Sum of formal charges = molecular or ionic charge
d Orbitals & Minimizing Formal Charge 2 O 6 + 4(6) + 2 = 32 valence electrons O S O 32/2 = 16 Shared/Lone Pairs O
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Better Lewis structure with minimized Formal Charge
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Note: There are resonance structures (draw these?) _____Formal Charges_____ S 6 – 6 = 0 O(single) 6 – 7 = -1 O(double) 6 – 6 = 0
More Formal Charge Examples H N H O C H H H N H H 1+ _____Formal Charges_____ C: O: N: H: 4 – 4 = 0 6 – 6 = 0 5 – 5 = 0 1 – 1 = 0 H: N: 1 – 1 = 0 5 – 4 = 1
Rules for Drawing Resonance Structures 2 2 O O O C C C O O O O O O 2 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 H 2 C O CH 3 vs.
H 2 C O 3. Charge Separation Reduces Stability CH 3 4. Negative Charge on More Electronegative Atoms
Isomerism: Structural
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Structural Isomers: Same Molecular Formula; Different Connectivity
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Why Might This Be a Big Deal? Consider Properties: C 2 H 6 O BP MP CH 3 CH 2 78.5 OH o C -117.3 o C CH 3 OCH -24.9 -138 o o C C 3
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Properties Can Differ Substantially Between Isomers!!
Cl Isomerism: Cis/Trans H H Cl C C C Cl H Cis or (Z) C Cl H Trans or (E)
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Same Molecular Formula (C 2 Cl 2 H 2 )
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Same Connectivity
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Different Structures
Double Bonds Don’t Rotate
Hybridization For now, worry only about Carbon hybridization Recall C’s valence configuration: 2s 2 2p 2 s orbital p orbital Will combine to form hybrid orbitals based on the valence of the carbon atom
Hybridization (2) Carbon Type Alkane Alkene Alkyne Hybridization sp 3 sp 2 (one pure p left) sp (two pure p left) 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. BF 3 , carbocations) Trigonal Pyramidal (e.g. NH 3 , carbanions) Tetrahedral (e.g. CH 4 , Ammonium Ion) Angular (Bent) (e.g. H 2 O)
Representations of Organic Structures
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Condensed Formula: CH 3 CH 2 OH, CH 3 CH 2 CH 2 CH 3
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Dash Formula: H H H H H H H C C H
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Bond-Line Formula H O H H C H C H C H C H H OH
Some Common Cyclic Structures H 2 C CH 2 H 2 C Cyclopropane H 2 C H 2 C CH 2 H 2 C H 2 C CH 2 C H 2 Cyclohexane CH 2 H 2 C CH 2 Cyclobutane HC H C H 2 C CH H 2 C HC CH C H Benzene CH 2 H 2 C CH 2 Cyclopentane