Chemistry 343—Summer 2006
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Transcript of Chemistry 343—Summer 2006
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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…
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Organic Chemistry: What and Why
• Compounds Based on Carbon
• Biological Molecules• DNA• RNA• Amino Acids/Proteins
• Photosynthesis
• Pharmaceuticals
• A #&*$ Load of Other Stuff
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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
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Empirical & Molecular Formula Examples
Consider 4 Hydrocarbons:
Ethene, Cyclopentane, Cyclohexane, 2-Butene
Empirical Formula: CH2
Molecular Formula: C2H4, C5H10, C6H12, C4H8
C C
H
HH
H
H3CCH3
H
H
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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)
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Valence and the Periodic Table
• Valence Corresponds To Column (Group I, II, Nonmetals)
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Electronegativity and the Periodic Table
• Know the electronegativity trends!!
Incr
easi
ng E
lect
rone
gativ
ity
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Lewis Structures
• Use only valence (outer shell) electrons
• Each atom acquires Noble gas configuration
• Octet Rule exceptions: Ions, Radicals, 3rd rowand lower (S, P, etc.)
• Sum # of valence electrons in atoms: this is the number of electrons that should be representedin the Lewis structure
• ½(valence electrons) = # shared + lone pairs
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Example: CH3Br
C
H
H Br
H
4 + 3(1) + 7 = 14 valence electrons
14/2 = 7 Shared/Lone pairs
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Example: C2H4
2(4) + 4(1) = 12 valence electrons
12/2 = 6 Shared/Lone Pairs
C C
H
H
H
H
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Example: CO32-
C
O
O O
2-
4 + 3(6) + 2 = 24 valence electrons
24/2 = 12 Shared/Lone pairs
• Place brackets around ions, indicate their charge
• We could have just as easily placed the double bond at other 2 O’s
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Resonance: The Carbonate Ion
C
O
O O
2-
CO
O
O
2-
CO O
O
2-
• Double headed arrows indicate resonance forms
• Red “Curved Arrows” show electron movement
• Curved Arrow notation used to show electron flow in resonancestructures as well as in chemical reactions: we will usethis electron bookkeeping notation throughout the course
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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
S
O
O
O
O
2-
6 + 4(6) + 2 = 32 valence electrons
32/2 = 16 Shared/Lone Pairs
• THIS IS NOT THE BEST POSSIBLE LEWIS STRUCTURE!
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Formal Charge• Formal Charge = #Valence Electrons - #Assigned Electrons
• We assign all electrons in a lone pair to an atom;½ bonded electrons
S
O
O
O
O
2-
S: 6 – 4 = +2
O: 6 – 7 = -1
Formal Charges
• Lewis structures that minimize formal charge tend to be better
• Note: Sum of formal charges = molecular or ionic charge
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d Orbitals & Minimizing Formal Charge
S
O
O
O
O
2-
6 + 4(6) + 2 = 32 valence electrons
32/2 = 16 Shared/Lone Pairs
S 6 – 6 = 0
O(single) 6 – 7 = -1
O(double) 6 – 6 = 0
_____Formal Charges_____ • Better Lewis structure with minimized Formal Charge
• Note: There are resonance structures (draw these?)
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More Formal Charge Examples
C
O
N HH
H
N
H
H H
H
1+
_____Formal Charges_____
C: 4 – 4 = 0O: 6 – 6 = 0N: 5 – 5 = 0H: 1 – 1 = 0
H: 1 – 1 = 0N: 5 – 4 = 1
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Rules for Drawing Resonance Structures
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
C
O
O O
2-
CO
O
O
2-
CO O
O
2-
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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
3. Charge Separation Reduces Stability
4. Negative Charge on More Electronegative Atoms
O CH3H2C vs. O CH3H2C
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Isomerism: Structural
• Structural Isomers: Same Molecular Formula; DifferentConnectivity
• 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!!
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Isomerism: Cis/Trans
C
C
Cl H
Cl H
C
C
H Cl
Cl H
Cis or (Z) Trans or (E)
• Same Molecular Formula (C2Cl2H2)
• Same Connectivity
• Different Structures Double Bonds Don’t Rotate
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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 valenceof the carbon atom
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Hybridization (2)
Carbon Type Hybridization Hybrid Composition
Geometry
Alkane sp3 25% s75% p
Tetrahedral
Alkene sp2
(one pure p left)33% s67% p
Trigonal planar
Alkyne sp(two pure p left)
50% s50%p
Linear
Hybrid orbitals form single () bonds; pure p form multiple ()
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VSEPR Theory: What to Know
You are responsible for these geometries (the mostprevalent 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)
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Representations of Organic Structures
• Condensed Formula: CH3CH2OH, CH3CH2CH2CH3
• Dash Formula:
• Bond-Line Formula
H C
H
H
C
H
H
O H H C
H
H
C
H
H
C C
H
H
H
H
H
OH
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Some Common Cyclic Structures
H2C CH2
CH2
H2C
H2C
HC
HCCH
CH
CH
HC
H2C
H2C CH2
CH2H2C
H2CCH2
H2C
H2CCH2
CH2
CH2
H2C
Cyclopropane Cyclobutane Cyclopentane
Cyclohexane Benzene