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1 H 3 C CH 3 H H OH C H H 3 C F HO C H H 3 C F O C O H 3 C H O H O C O H 3 C O H H + + X H H 2 / M Cl I C H H 3 C D NC C H H 3 C D (-) :CN: ( - ) I CH 2 OH O H CH 3 CH 3 H Br Mg CH 3 LiCu 2

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1

H3C

CH3

H

H OH

CH

H3C F

HO

CH

H3CF

O

CO

H3C

H

OH

O

CO

H3C

OH

H

+

+

XH

H2 / M

Cl

IC

H

H3C D

NC

CH

H3CD

(-):CN:(-)I

CH2

OH O

HCH3

CH3

H

Br Mg CH3

LiCu

2

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2

Identify a seat (seating chart will be circulated)Lecture Recitations start this week

- work problems, discuss concepts, take quizzesPrerequisites and other things

Start Chapter 1

1st labs meet this week Non-major’s lab 331L

Major’s lab 333L

Contact information

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3

Professor John H. Dawson, GSRC Room 410, 7-7234, [email protected]

Lecture: TuTh 5:00-6:15 PM PSC Room 002

Office Hours: M-Th 3:00-3:45 PM, before/after class and by appointment

Lecture Recitation Sections: Section 1, Room PSC 115, Fri 1:25-2:15 PM, Vibha Gupta Sect 2, Room PSC 203, Thurs 12:30-1:20 PM, Lisa Brodhacker Sect 3, Room PSC 214, Fri 1:25-2:15 PM, Avneesh Saini

Sect 4, Room PSC 214, Thurs 2:00-2:50 PM, Andrew Lee

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4

choose your “assigned seat”

FrontDoor

front desk

screen-white boards

PSC002BackDoor

1 2 3 4 5 6 7 8 9 10 11 12 13 14

J I

H G FEDCBA

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Prerequisites and other Things

You must have passed Chem 112 with a C/D (major/non-major) or better to be in Chem 333.

You must be registered for the course unless you are making up an incomplete (I).

If you are making up an incomplete, see JHD

Grades will be posted by the last four numbers of your SSN

Or

Give me an “alias” set of four numbers to use in place of the last four numbers of your SSN

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6

Text: Organic Chemistry, 4th ed. Brown, Foote, Iverson (3rd ed.?)

Also: Study Guide and Problems & models are strongly recommended!

Consider other aids: Organic As a Second Language OWLDifferent organic textsOthers?

Exams: Jan 31, Mar 2, Apr 11Final: May 1, 5:30 PM

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7

1.1 Electronic Structure of Atoms1.2 Lewis Model (Octet Rule, Formal Charge) 1.3 Functional Groups 1.4 Bonding Angles and Shapes of Molecules 1.5 Polar and Nonpolar Molecules1.6 Resonance1.7 Quantum or Wave Mechanics1.8 Molecular Orbital & Valence Bond Theory, Covalent

Bonds, Hybridization (sp3, sp2, sp) SUMMARY and OVERVIEW 1-11

COVALENT BONDING & SHAPES OF MOLECULES

1

Chapter 1

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8

Li Be B C N O F Ne WHY is CARBON SPECIAL?

Organic Chemistry: the study of compounds of carbon

CARBON - is small, intermediate electronegativity - forms strong bonds with itself/other atoms C-C 83.1 kcal/mole

2X as strong as: N-N, O-O, Si-Si

Over 10 million structures identified; ~1000 new/day!

-forms strong double and triple bonds (to C or other atoms)

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9Recall the Structure of an ATOM:

1833 to 1

electrons(-)

nucleus(neutrons + protons(+))

mass:

quantum mechanics: motion of e’s particle & wave like

electrons confined to regions of space: shells (principle energy levels)

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each shell can hold 2n2 electrons

Electronic Structure of Atoms

Shell

e-s Shell

can hold

Relative Eenergy

of e-s in Shells

321

18 8 2

higher

lower

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11

ground state electronic configuration(atoms or molecules)

Aufbau Principle:Aufbau Principle:

fill orbitals lowest to highest energy

Pauli Exclusion Principle:Pauli Exclusion Principle:

2 electrons per orbital, spins paired

Hund’s Rule:Hund’s Rule:

degenerate orbitals, 1 electron in each

then create a pair

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12

Hund’sRule

next level

H He

Li Be B C N O F Ne

Valence Shell

E

1

2

3

Pauli&

Lewis

(s)

(s)

(p)

(s)

closed shell

writing electronic configuration

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13

No. of e's electronic configuration1 H 1s1

2 He 1s2

3 Li 1s2 2s1

4 Be ……. 5 B ………6 C 1s2 2s2 2p2

7 N etc.

Electronic Configuration

What is the electronic configuration of OXYGEN?

1s2, 2s2, 2p4

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14

Lewis Structures (Gilbert N. Lewis)

Valence shell:Valence shell: the outermost electron shell

Valence electrons:Valence electrons:

electrons in valence shell

electrons used to bonds

Lewis structure:Lewis structure: atom symbol = nucleus + inner e-

dots represent valence electrons

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15

Lewis Structures• Table 1.4 Lewis Structures

H

Li

Na

1A 2A 3A 4A 5A 6A 7A 8A

.

.

.

Be

Mg

:

:

C

Si

.

.

:

:.

. N O

Cl

F

SP ::

..

..

. .

.

.

.::::

:::::.

:.

.

He

Ne

Ar

::

::

:

:

:

:

:

B

Al

.

.

:

:

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electrons to complete a shell

H He+1 0Li Be B C N O F Ne-1 -2 -3 ±4 +3 +2 +1 0

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17

Bonding extremes - IONIC or COVALENT

ionic bonds ?

Li Br

- loss or gain of valence electrons

Li+ Br

covalent bonds ?

Br Br

-share electrons to fill shells

Br Br Br Bror

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ElectronegativityElectronegativity:Electronegativity: a measure of an atom’s

attraction for the electrons it shares with another atom

Pauling scaleincreases left to right in a rowincreases bottom to top in a column

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19

ElectronegativityTable 1.6 Classification of Bonds

electronegativity difference bond type

H3C-H 2.5 2.1

less than 0.5 covalent

greater than 1.9 ionicNa+ -Cl

0.93 3.16

0.5 to 1.9 polar covalentH Cl2.1 3.0

+ -

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1.1 Electronic Structure of Atoms1.2 Lewis Model (Octet Rule, Formal Charge) 1.3 Functional Groups 1.4 Bonding Angles and Shapes of Molecules 1.5 Polar and Nonpolar Molecules1.6 Resonance1.6 Quantum or Wave Mechanics1.7 Molecular Orbital & Valence Bond Theory, Covalent

Bonds, Hybridization (sp3, sp2, sp) SUMMARY and OVERVIEW 1-11

COVALENT BONDING & SHAPES OF MOLECULES

1

Chapter 1

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21

Lewis (electron dot) structures covalent molecules and ions

Determine the number of valence e's in molecule (ion) Ions - add/subtract 1 electron for -/+ charge

Determine an order of attachment

Connect atoms with single bonds. Arrange remaining e's in pairs so each atom has a complete shell.

Pairs of e's - shown as “ - “; nonbonding e's as dots

Shared e's can be single, double or triple bonds

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22

Charge on an atom/molecule is formal charge, i.e. H3O

+, HO-

To assign Formal Charge:1. Write correct Lewis structure2. Assign each atom: all non-bonding e's

half the shared e’s

3. Compare this number to valence e-s in neutral unbonded atom.

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23

Formal Charge on an atom

Formal Charge

unshared electrons +

1/2 of shared electrons

-number of valence

electrons

=

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valence of N = 5

e’s belong to the first N =

NNNH

e’s belong to the 2nd N =

e’s belong to the 3rd N =

[2 + 1/2(6)] = 5

5 - 5 = 0 charge

[1/2(8)] = 4

5 - 4 = +1 charge

[4 + 1/2(4)] = 6

5 - 6 = -1 charge

0 +1 -1

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25

Exceptions to the octet rule:

Group 3A - B, Al - due to valence, form 3 bonds - stable but reactive

3rd row+ elements - e.g. S - can share more than 8 e’s

H3CS

O

CH3

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H N C:

Recall Formal Charge

Lewis structures - filled shells

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1.1 Electronic Structure of Atoms1.2 Lewis Model (Octet Rule, Formal Charge) 1.3 Functional Groups 1.4 Bonding Angles and Shapes of Molecules 1.5 Polar and Nonpolar Molecules1.6 Resonance1.6 Quantum or Wave Mechanics1.7 Molecular Orbital & Valence Bond Theory, Covalent

Bonds, Hybridization (sp3, sp2, sp) SUMMARY and OVERVIEW 1-11

COVALENT BONDING & SHAPES OF MOLECULES

Chapter 1

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Determine:reactionsproperties basis of nomenclature & classification

Functional Groups: atom(s) bonded to C having characteristic properties

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alcohols: hydroxyl group C OH

amines:amino group C NH

H

here

methyl,1o, 2o, 3o

see Hs & Cs

1o, 2o or 3o by Hs on N

Functional Groups: atom(s) bonded to C having characteristic properties

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FUNCTIONAL GROUPS

alcohols: (hydroxyl group, sp3-> C-O-H)

Formulas: complete, condensed and/or line

Aldehyde / Ketone (carbonyl group C=O)

Carboxylic Acid (carbonyl + hydroxyl group -CO2H)

Carboxylic Ester - (carbonyl + alcohol -CO2R)

amines: 1o, 2o, 3o

1o, 2o, 3o

Others ethers, halides, etc.

Carboxylic Amide - (carbonyl + amine -C(O)N-)

C N H

H

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VSEPR - electrons (e-) in bonds/orbitals repel each other

4 groups/bonds repulsion yields a tetrahedral shape ~109.5o

3 bonds repulsion yields a trigonal shape ~120o

2 bonds repulsion yields a linear shape 180o

eg: CH4 or NH3, HCO2H, CO2

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and Nonpolar Molecules

Polar if:

(2) if the arrangement is “irregular”

(1) has polar bonds

O C O

H N C:+ -

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Resonance - sometimes a Lewis formula is not an accurate representation.

H C

O

N H

e.g. the anion of formamide:

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- sometimes a Lewis formula is not an accurate representation.

Resonance

H C

O

N H H C

O

NH

note: charges and double bonds

Compound w/ >1 Lewis structure but same connectivity = resonance

H C

O

NH

Resonance hybrid- real, average

- connect contributors with

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Create new contributing structures by arrows - e's move from source to need!

1. Flow from an atom to an adjacent bond

2. Flow from a bond to an adjacent atom

source:

non-bonding, -bond, p-orbital e's

Resonance - some Lewis formulas inaccurate representation

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37Writing acceptable resonance contributing structures

1. Contributing structures-same # valence e-s.

2. 2nd row elements - no more than 8 e-s, 3rd row up to 12e-s

3. Nuclei don’t change positions4. All contributing structures should show

the same number of paired and unpaired electrons

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1. Same number of valence e-s

Writing acceptable resonance contributing structures

H C

O

O

6

4

7

H C

O

O 7

4

6

2. 2nd row 8; 3rd row > 8

H C

O

O

H C

O

O

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valid

3. All nuclei remain in the same positions

Rules to write acceptable resonance contributing structures

H C

O

O

H C

O

OH

C

O

O

not valid!

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4. Same number of paired & unpaired e-s.

not valid e-pairs have decreased

Rules to write acceptable resonance contributing structures

H2C CH2 H2C CH2X

H2C CH2 H2C CH2this is “OK”

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41

Which resonance contributing structure(s) is(are) IMPORTANT

1. complete octets contribute more.

2. more covalent bonds - contribute more.

3. charge separation less important.

4. (-) on more electronegative atom, (+) on electropositive is more important.

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42

http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/FG12_05.JPG

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Orbitals for sp2 hybridization

3 sp2 and p orbitals on axis

hybrid

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47

3 sp2 and p orbitals on axis

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49pz

py

s

px

two sp hybrid orbitals

pz

sptwo sp hybrid orbitals

py

sp

sp HYBRIDIZATION

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End of Chapter 1

2