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

Empirical Formula: CH2

Molecular Formula: C2H4, C5H10, C6H12, C4H8

C C

H

HH

H

H3CCH3

H

H

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)

Electronegativity and the Periodic Table

• Know the electronegativity trends!!

Incr

easi

ng

Ele

ctro

neg

ativ

ity

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

Example: CH3Br

C

H

H Br

H

4 + 3(1) + 7 = 14 valence electrons

14/2 = 7 Shared/Lone pairs

Example: C2H4

2(4) + 4(1) = 12 valence electrons

12/2 = 6 Shared/Lone Pairs

C C

H

H

H

H

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

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

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!

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

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?)

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

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-

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

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!!

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

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

Hybridization (2)

Carbon Type Hybridization Hybrid Composition

Geometry

Alkane sp3 25% s

75% p

Tetrahedral

Alkene sp2

(one pure p left)

33% s

67% p

Trigonal planar

Alkyne sp

(two pure p left)

50% s

50%p

Linear

Hybrid orbitals form single () bonds; pure p form multiple ()

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)

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

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