Chapter 1Chapter 1Introduction and ReviewIntroduction and Review

Organic Chemistry, 6th EditionL. G. Wade, Jr.

Chapter 1 2

DefinitionsDefinitions

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

heatNH4

+ OCN- H2N C NH2

O

urea

Chapter 1 3

Atomic StructureAtomic 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

Chapter 1 4

Electronic StructureElectronic Structure

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

the electron in a particular part of an orbital.

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

Chapter 1 5

First Electron Shell

The 1s orbital holds two electrons.

Spherical shape

Chapter 1 6

Second Electron Shell

Three p orbitals

P-orbital (dumbell shaped)S-Orbital (spherical)

Second shell (4 orbitals) holds total of 8 Second shell (4 orbitals) holds total of 8 electronselectrons

Chapter 1 7

Electronic ConfigurationsElectronic Configurations• Aufbau principle:

Place electrons in lowest energy orbital first.

• Hund’s rule: Equal energy orbitals are half-filled, then filled.

• 6C: 1s2 2s2 2p2

Chapter 1 8

Electronic Configurations

=>

Chapter 1 9

Bond FormationBond Formation

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

Chapter 1 10

Lewis StructuresLewis Structures• Bonding electrons represented by a dash line

or a pair of electron• Nonbonding electrons or lone pairs are shown

Satisfy the octet rule!

C

H

H

H

OH

C :....: O:HH

H

H

..

..

Chapter 1 11

Multiple Bonding

=>

Chapter 1 12

Dipole MomentDipole MomentAmount of electrical charge x bond length.

Chapter 1 13

Electronegativity and Bond PolarityElectronegativity and Bond Polarity

Greater EN means greater polarity

=>

Chapter 1 14

Calculating Formal ChargeCalculating Formal ChargeFor each atom in a valid Lewis structure:• Count the number of valence electrons• Subtract all its nonbonding electrons• Subtract half of its bonding electrons

C

H

H

H

C

O

O P

O

OO

O

3-

F.C = # valence e – (non-bonding e + 1/2(bonding e)

Chapter 1 15

Ionic StructuresIonic Structures

C

H

H

H N

H

HH

+

Cl-

Na O CH3 or O CH3Na+ _X

=>

If the EN is greater than 2.5, the bond is a ionic bond.

Chapter 1 16

ResonanceResonance• Only electrons can be moved (usually

lone pairs or pi electrons).• Nuclei positions and bond angles remain

the same.• The number of unpaired electrons

remains the same.• Resonance causes a delocalization of

electrical charge.

Chapter 1 17

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

Chapter 1 18

Major Resonance Form or Major Resonance Form or Major ContributorMajor Contributor

• has as many octets as possible.• has as many bonds as possible.• has the negative charge on the most

electronegative atom.• has as little charge separation as

possible.

Chapter 1 19

Resonance Hybrid

majorcontributor

minor contributor, carbon does

not have octet=>

Chapter 1 20

Chemical FormulasChemical Formulas

• Full structural formula (no lone pairs shown)

• Line-angle formula

• Condensed structural formula

• Molecular formula• Empirical formula

CH3COOH

C2H4O2

CH2O

C

H

H

H

C

O

O H

OH

O

Chapter 1 21

Calculating Empirical FormulasCalculating Empirical Formulas

• Given % composition for each element, assume 100 grams.

• Convert the grams of each element to moles.

• Divide by the smallest moles to get ratio.• Molecular formula may be a multiple of

the empirical formula.

Chapter 1 22

Sample ProblemAn unknown compound has the following composition: 40.0% C, 6.67% H, and 53.3% O. Find the empirical formula.

molCmolCgCgC 33.3/0.12

0.40

molHmolHgHgH 60.6/01.1

67.6

molOmolOgOgO 33.3/0.16

5.53

3.33

3.33

3.33

= 1

= 1

= 1.98 = 2

Empirical Formula: CH2O

Chapter 1 23

Arrhenius Acids and BasesArrhenius Acids and Bases

• Acids dissociate in water to give H3O+ ions.• 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.

1 M1 MCl

-+H3O

+H2O+HCl

Chapter 1 24

BrBrØØnsted-Lowry Acids and Basesnsted-Lowry Acids and Bases

• Acid: donates a proton.• Base: accepts a proton.• Conjugate acid-base pairs.

CH3 C

O

OH + CH3 NH2 CH3 C

O

O- + CH3 NH3+

acid conjugatebase

base conjugateacid

=>

Chapter 1 25

Acid and Base StrengthAcid and Base Strength

• Acid dissociation constant, Ka

• Base dissociation constant, Kb

• For conjugate pairs, (Ka)(Kb) = Kw

• 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

Chapter 1 26

Structural Effects on AcidityStructural Effects on Acidity

• Electronegativity • Size • Resonance stabilization of conjugate

base

Chapter 1 27

ElectronegativityElectronegativityAs the bond to H becomes more polarized, H becomes more positive and the bond is easier to break.

=>

Chapter 1 28

SizeSize• As size increases, the H is more loosely

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

Chapter 1 29

ResonanceResonance• Delocalization of electron pairs.• More resonance structures usually

mean greater stabilization.

CH3CH2OH < CH3C

O

OH < CH3 S

O

O

OHAcidity:

Chapter 1 30

Lewis Acids and BasesLewis Acids and Bases• Acids accept electron pairs = electrophile• Bases donate electron pairs = nucleophile

=>

Chapter 1 31

End of Chapter 1