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1

Chapter 1Introduction and Review

Chemistry 231Organic Chemistry I

Fall 2007

Introduction and Review Slide 1-2

Definitions• Old: “derived from living organisms”• New: “chemistry of carbon compounds”• From inorganic to organic, Wöhler, 1828

heatNH4

+ OCN

- H2N C NH2

O

urea


Atomic Structure• Atoms: protons, neutrons, and electrons.• The number of protons determines the identity of

the element.• Some atoms of the same element have a different

number of neutrons. These are called isotopes.• Example: 12C, 13C, and 14C• Isotopes of a given element, when incorporated

into molecules, behave no differently than anyother isotope of that element**:

12CO2, 13CO2, and 14CO2 all chemically identical**

2


Electronic Structure• Electrons: outside the nucleus, in orbitals.• Electrons have wave properties.• Electron density is the probability of finding

the electron in a particular partof an orbital.

• Orbitals are grouped into “shells,” atdifferent distances from the nucleus.


First Electron Shell

The 1s orbital holds two electrons.


Second Electron Shell

2s orbital (spherical)

Three p orbitals

2p orbital

3


Electronic Configurations

• Aufbau principle:Place electrons inlowest energy orbitalfirst.

• Hund’s rule: Equalenergy orbitals arehalf-filled, then filled.

• 6C: 1s2 2s2 2p2

↑↓

↑↓

↑ ↑


Electronic Configurations

=>


Bond Formation

• Ionic bonding: electrons are transferred.• Covalent bonding: electron pair is shared.

4


Lewis Structures• Bonding electrons• Nonbonding electrons or lone pairs

Satisfy the octet rule!

C

H

H

H

OH


Drawing Lewis Structures-11. Determine the total number of valence electrons for ALL

atoms.a) Don’t be concerned with which atom gave what; the SUM

TOTAL is what is important. If the entire molecule is charged(i.e. a polyatomic anion or cation) add one valence electron foreach unit of negative charge (if it is an anion) and remove onevalence electron for each unit of positive charge (if it is a cation).

2. Write a skeleton structure for the molecule, making theleast electronegative atom the central atom.a) The order of electronegativity for the nonmetals is

F>O>N>Cl>Br>I>S>C>H. Hydrogen is never the central atom.b) If there are more than one of the least electronegative atom, your

skeletal structure should have those two attached to one another(EXCEPTION: Hydrogen)


Drawing Lewis Structures-23. Connect each member of the skeleton structure to the central atom(s) using a

single line to represent a bond.Each bond is comprised of two electrons so each line indicates a two-electronbond.

4. Determine the number of electrons you have left to distribute.

5. Starting with the atoms bonded TO the central atom (the ‘outside’ atoms),distribute the electrons two at a time until all electrons from step 4 are usedOR until each atom has an octet (eight electrons around it); DO NOT givehydrogen any of these ‘left-over’ electrons.

A good rule of thumb is to give EACH non-hydrogen two electrons at a time untileach outside atom has two extras, then give two more to each non-hydrogen untileach has four extra electrons and so on. Once all of the ‘outside’ atoms haveoctets (except for the hydrogens), put remaining electrons on the central atom(s)until all electrons are used or every atom has an octet (or a duet, in the case of H).

AT THIS POINT: Anytime you run out of electrons, look at each atom and determine ifit does or does not have an octet. If all atoms have an octet go on to Step 8. If not,proceed to Step 6.

5


Drawing Lewis Structures-36. If there are too few electrons available to complete octets for all the atoms

that need them, make double and/or triple bonds between appropriate atoms.• Remember that in doing this, no atom should lose electrons (i.e. double and triplebonds SHARE electrons BETWEEN the two involved atoms and as a resultneither atom involved in the multiple bond loses electrons, so do not moveelectrons FROM one atom to another).

• A major exception to this rule is one regarding electrically-neutral Lewis acidscontaining Be, Al or B: it is possible to find Be with 4e- around it (BeCl2), and Band Al with 6e- around them (AlCl3 and BF3, for example)

7. If there are too many electrons available: first, RECOUNT the valenceelectrons available; then, after completing octets for all the atoms that needthem, place remaining electrons on the central atom IF the central atom is aPeriod 3 or greater element.

8. After you have completed your Lewis structure, CHECK FOR OCTETS. Ifall atoms have an octet, calculate formal charges for all atoms.


Drawing Lewis Structures-Formal ChargeCalculation of formal charge, while not difficult, is the most common mistake

made by students at all levels of chemistry when drawing Lewis structures.Even neutral molecules MAY have individual atoms that bear a formalcharge Formal charge is calculated for each atom independent of any othersin the molecule. Remember: if you are drawing the Lewis structure of acation or an anion, it WILL have a charge (or charges) in it.

The procedure is as follows:1) Determine the number of valence electrons (VE) that the atom of

interest has; this is most easily done by using a periodic table;2) Determine the number of unshared (or non-bonded) electrons (NBE)

that the atom of interest has;3) Determine the number of bonds (B) that the atom of interest has

(remember that a double bond counts as two and a triple bond counts asthree);

4) Formal Charge is then calculated using the following formula:FC = VE - B - NBE


Lewis Structures-Formal Charge Example

Let’s calculate the formal charge for all the atoms in the Lewisstructure for SO2 shown below:

For Oxygen 1: FC = 6 - 1 - 6 = -1;For Sulfur: FC = 6 - 3 - 2 = +1;

For Oxygen 2: FC = 6 - 2 - 4 = 0;

Thus the complete structure looks like the following:

S OO

1 2

S OO

6


Formal Charge ExamplesWhat are the formal charges of the indicated atoms?

C ClCl C ClCl

H

C OO

C HH

C HH

C

C

N

C

C

C

CH3

H

H

H

H H

C ClCl

H


Formal Charge Answers

C ClCl C ClCl

H

C OO

C HH

C HH

C

C

N

C

C

C

CH3

H

H

H

H H

C ClCl

H0 0

0 0

+1

+1

-1

-10

a carbene a radical a carbanion

a carbocation

(carbonium ion)


Estimating “Correctness” of LewisStructures

What do you do if more than one Lewis structure can be drawnfor a given molecule?

1. More covalent bonds, more stable and therefore better thestructure;

2. Charge separation will decrease stability (smallest amountof charge separation = better structure);

3. Those structures where all atoms have an octet will be thebest ones (more octets = better structure);

4. Minimization of charge is best, with negative charge onthe most electronegative atom.

7


Lewis Structure: Final Notes• Organic structures are typically drawn in one of three ways:

Extended Lewis, Condensed Lewis or Line-Angle; a mix ofall three is common (more in lab lecture).

• Lone-pairs are typically not shown unless discussingreaction mechanisms; they are understood to be present.

• Formal charges of zero are never shown.• Resonance will be an important adjunct to Lewis structures.• 3-Dimensional structure is also very important (later).• YOU MUST BE ABLE TO RAPIDLY DRAW AND

INTERPRET LEWIS STRUCTURES TO BESUCCESSFUL IN ORGANIC!


Dipole Moment and Polar Molecules• Amount of electrical charge × bond length.• Charge separation shown by electrostatic potential map

(EPM).• Polarity is a function of 3D structure!• Red indicates a partially negative region (δ-) and blue

indicates a partially positive region (δ+).


Electronegativity and Bond PolarityGreater ΔEN means greater polarity

C-N bond ‘more’ covalentthan C-F bond

Differences in electronegativity lead to polarbonds!

8


Ionic Structures

C

H

H

H N

H

H

H

+

Cl-

Na O CH3 or O CH3Na+_X

Organic molecules can be ionic as well:

Note: this representation is NOTincorrect, simply misleading


Resonance

Resonance isomers: isomers that have the same sigmaconnections of atoms but differ in the positions of theremaining electrons. Resonance isomers of a givenmolecule MUST be interconvertable with each otherAND MUST be interconvertable by ONLY movingelectron pairs.

• Only electrons can be moved (usually lone pairs or pielectrons).

• Nuclei positions and bond angles remain the same.• The number of unpaired electrons remains the same.• Resonance causes a delocalization of electrical charge.


Resonance IsomersWhen drawing Lewis structures of molecules that have unshared pairs ofelectrons and/or multiple bonds, occassionally more than one structurecan be drawn differing only in the location of electron pairs. ConsiderCO2:

Both Structure a and b are valid Lewis structures and areinterconvertable by moving pairs of electrons:

These two structures are are called resonance structures or resonanceisomers; the double headed arrow (↔) is called a resonance arrow.BOTH are correct; but which one is better?

O C O O C O

Structure a Structure b

O C O O C O

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Another Resonance Example

• The real structure is a resonance hybrid.• All the bond lengths are the same.• Each oxygen has a -1/3 electrical charge.

N

O

OO

_ _

N

O

OO

_

N

O

OO


Resonance Guidelines-11. Resonance structures exist only on paper. Resonance is a

theory that allows us to describe molecules, ions, or radicalsfor which a single Lewis structure is inadequate. Resonanceallows one to envoke the idea of electron and chargedelocalization and helps to explain chemical differences thatare almost impossible to explain otherwise.

2. When writing resonance structures only electrons are allowedto be moved. This is to say that for any set of resonancestructures, the sigma (σ) network can NOT change and thepositions of the nuclei of the atoms must remain the same inall structures. This means that in general we move ONLYelectrons associated with π bonds and unshared electron pairs.


Resonance Guidelines-23. All resonance structures for a given compound must be proper

Lewis structures. The rules of valency must be obeyed; allresonance structures must have the same total number ofelectrons. It is OK (but not ideal) for one of a pair ofresonance structures to have an incomplete valence (openoctet).

4. All resonance structures must have the same number ofunpaired electrons. Breaking a double bond into two unpairedelectrons does not constitute a resonance isomer.

5. All atoms that are involved in resonance must lie (or be ableto lie) in the same plane. This is due to the fact that for orbitalsto interact with one another they have to be coplanar.

10


Resonance Guidelines-36. The overall energy of the actual molecule will be lower than

any given contributing structure. This is resonancestabilization and is due to the fact that the actual molecule is ahybrid of all possible contributing resonance structures.

7. The more equivalent resonance structures possible for a givenmolecule, the larger the resonance stabilization energy willbe. Equivalent resonance structures make equal contributionsto the hybrid so the more of them you can have the more‘spread out’ or delocalized the electrons will be and the morestable the molecule will be.

8. The more stable a resonance structure is, when evaluated onits own, the greater its contribution is to the hybrid; theopposite is also true. Since the hybrid is the weighted sum ofall the contributors, the more stable a contributor is the moreimportant its contribution.


Resonance Illustrations

O O

C

O

O

O

CO

O

O

CO

O

O

C

O

-2/3

-2/3

-2/3

Hybrid

O O

C

O

OK

O O

C

O

Not OK


Resonance Hybrid

majorcontributor

minor contributor, carbon does

not have octet

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Structural Representations• Full structural formula (no lone pairs shown)

• Line-angle formula

• Condensed structuralformula

• Molecular formula

• Empirical formula

CH3COOH

C2H4O2

CH2O

C

H

H

H

C

O

O H


Acid-Base Chemistry• One of the most important underlying principles in

understanding organic reactions;• Three distinct definitions employed:

ArrheniusBronsted/LowryLewis

• Predictions of reaction outcome possible using basicacid-base principles.


Arrhenius Acids and Bases• Acids dissociate in water to give H3O+ ions (“H+”).• Bases dissociate in water to give OH- ions.• Kw = [H3O+ ][OH- ] = 1.0 x 10-14 at 24°C• pH = -log [H3O+ ]• Strong acids and bases are 100% dissociated.

• Weak acids and bases dissociate to only 5-10%

HCl

1 M

! "# + H

2O ! "!

# !! H3O

+

1 M

! "##+ Cl

–

CH3CO

2H

1 M

! "###### + H

2O ! "!

# !! H3O

+

>>1 M

! "##+ CH

3CO

2

–

12


Bronsted-Lowry Acids and Bases

• Acids and bases defines in their ability totransfer or accept hydrogen ions (“protons”)

• Acids can donate a “proton” (“H+”).• Bases can accept a “proton”.• Equilibrium control important.• Conjugate acid-base pairs.

CH3 C

O

OH + CH3 NH2 CH3 C

O

O-+ CH3 NH3

+

acid conjugatebase

base conjugateacid


Acid and Base Strength• Acid dissociation constant, Ka (very important)• pKa = -log Ka

• Base dissociation constant, Kb

• For conjugate pairs, (Ka)(Kb) = Kw

• pKa + pKb = 14• Spontaneous acid-base reactions proceed from stronger to

weaker.

CH3 C

O

OH + CH3 NH2 CH3 C

O

O-+ CH3 NH3

+

pKa 4.74 pKb 3.36 pKb 9.26 pKa 10.64


pKa Tables• Since conjugates are related, easier to tabulate Ka or pKa

values and use these to predict reactions;• Rules of thumb:

smaller the Ka, the less acidic the acid; larger the Ka, the more acidic the acid; smaller the pKa, the more acidic the acid; larger the pKa, the less acidic the acid;

• Determine the acid on both the left and the right side;• Estimate the pKa for each; reaction will proceed in the

direction that goes from the strongest acid to the weakestacid.

• pKrxn = pKa(left side) + pKa(right side)

13


pKa Tables


Structural Effectson Acidity

• Electronegativity• Size• Resonance stabilization of conjugate base


Electronegativity

As the bond to H becomes more polarized, Hbecomes more positive and the bond is easierto break.

14


Size (Bond Length)• As size increases, the H is more loosely held and

the bond is easier to break.• A larger size also stabilizes the anion

(delocalization).


Resonance• Delocalization of the negative charge on the

conjugate base will stabilize the anion, so thesubstance is a stronger acid.

• More GOOD resonance structures usuallymean greater stabilization.

CH3CH2OH < CH3C

O

OH < CH3 S

O

O

OH


Lewis Acids and Bases• Acids accept electron pairs = electrophile• Bases donate electron pairs = nucleophileThis relationship will be very important to keep in mind during

all of organic chemistry

15


Curved Arrow Formalism• Mechanism is important in organic chemistry; this is

typically indicated using curved arrow notation:TWO electrons flow

FROM here TO here FROM here TO here

ONE electron flows

• Electron flow is always FROM high electron density TO lowelectron density (- or δ- TO + or δ+); multiple arrows can bestrung toegether:

O O

C

O

H+ NotO O

C

O

H+

HO

H

C

H

H3C

H

Cl

CH3C HOH C C

H3C

CH3

H

H

Cl-+ +


Chapter 1 Homework:

20, 22, 24-29, 31, 34, 36, 37, 40, 44, 45

pKa Table (H2O Reference)

Oxygen Acids

CF3SO2H -14

N

O

O-HC6H5

-12.4

CO-H

O-HC6H5

-7.8

O

H

CH3C6H5

-6.5

CO-H

CH3C6H5

-6.2

(CH3)2S+—H -5.4

S H

-4.4

(CH3)2O+—H -3.8 CH3SO2-OH -2.6

CH3OH2+ -2.2

O H

-2.05

(CH3)2S=O+—H -1.8

N

O

O-HO

-1.37

CF3CO2H -0.25 NO2

NO2

O2N OH

0.3

C6H5-CO2H 4.2 CH3-CO2H 4.76

C6F5-OH 5.5

CO

O-HH-O

6.3

C6H5-SH 6.5 C6H5-OH 9.95

HCO3- 10.3 R-S-H 10.5 HO-H 15.7

CH3-OH 16 (CH3)3C-OH

[t-BuOH]

20

Nitrogen Acids

+PH4 -14

C6H5-C≡N+—H

-10.5

CH3-C≡N+—H

-10 NO2

NO2

O2N NH3

-9.3

N N

H

-2.9

(C6H5)2NH2+ 0.78

CH3-PH3+ 2.7

C6H5-NH3+ 4.6

N H

5.21

N≡C-CH2-CH2-NH3+ 7.87

(CH3CH2)3-PH+ 9.1

N HH2N

9.2

+NH4 9.2

CH3CH2-NH3+ 10.6

(CH3CH2)2-NH2+ 11.0

(CH3CH2)3-NH+ 10.75

CNH2

N(CH3)2(CH3)2N

13.6

C6H5-NH2 28 NH3 33

Carbon Acids

(NO2)2CH2 3.6

H

O

O

H

5.2

O O

H H

9

N≡C-H 9.1 O2N-CH3 10.2

OEt

O O

H H

10.7

(CF3)3CH 11

(N≡C)2CH2 11.2

EtO OEt

O O

H H

13

H

H

15

H3C C CH3

O

20

H H

20

CH2-CN

20.8

C6H5-C≡C-H 23

CH3-CO2Et 24 CH3-SO2-CH3 33

CH3-C≡N 25

H-C≡C-H 24 H

H

29

CF3-H 32 (C6H5)3CH 31.5

H3C S CH3

O

33.5

C6H5-CH3 41 CH2=CH2 44

H

43

H

H

46

CH3-CH3 50

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