PART 1

Introduction to Organic Chemistry

Chapter 1. Structure, bonding, acid-base2. Nomenclature, physical properties,

drawing structure

Chapter 1

Structure and BondingAcids and Bases

Electronic structureChemical bondsAcid-base

vital force theory by Berzelius, 1807 compounds from living organisms ~ ‘organic’ comp’ds from minerals ~ ‘inorganic’

death of VFT synthesis of urea (‘organic’) from ‘inorganic’ Wöhler, 1828

the current definition of organic comp’ds “compounds that contain carbon”

Organic compounds Ch 1 #3

VFT is dead!! VFT-1: Organic compounds from living organism (plant or

animal) only

VFT-2: An organic compound still contains some of the life force of the organism that makes them. vitamin C, saponin, etc

natural vs synthetic (p3) Vitamin C is vitamin C!

vitamin C ≡ L-ascorbic acid

Ch 1 #4

Chemistry of carbon comp’d organic chemistry = chemistry of carbon comp’d

why carbon? forms stable covalent bond to other carbon sharing electrons Li+, F- vs C4+ or C4-

forms chain variety (16M org comp’ds)

forms bonds to heteroatoms (O, N, S, P, X, etc.)

exceptions: CO, CO2, Na2CO3, etc. ~ inorganic

Ch 1 #5

Components of org chem organic compounds structure molecule bonding atoms

properties physical

chemical

Chapter 1, 2, 3, 5, 7

organic (chemical) reactions mechanism

thermodynamics

kinetics

Chapter 4, 6, 8, 9, 10, 11

Ch 1 #6

In org chem 1…Chapter

1. electronic structure and bondingacid-base

2. nomenclature, physical propertiesstructure drawing, alkanes

3,4. alkenes

5. stereochemistry

6. alkynes

7. resonance structure

8-10. substitution and elimination reactions

11. organometallics

Ch 1 #7

Structure of an atomDefine the followings atomic number

mass number

isotope

allotrope

atomic weight

atomic mass

molecular weight [molar mass]

amu

Ch 1 #8

Electronic structure of an atom quantum mechanics (Schrödinger) wave equation – wave function [orbital]

quantum numbers – shell – sub-shell [orbital]

electronic configuration ~ distribution of e in orbital aufbau principle ~ ‘building-up’

Pauli exclusion principle ~ two e with opposite spin

Hund’s rule ~ degenerate orbitals

Ch 1 #9

Bonding core vs valence electron

octet rule by Lewis ~ ‘Lewis structure’

8 e at valence shell ~ stable atom

Atoms lose, gain, or share eto satisfy octet rule and form bonding. ~ ‘Lewis theory’

Ch 1 #10

Ionic bond atom ion bonding

ionic compound ionic solid high mp

strong bond + network

ionic molecule?

Ch 1 #11

Covalent bond sharing e

valency mono, di, tri, or tetravalent

covalent comp’d ‘molecule’

weak intermol interaction

Ch 1 #12

Polar (covalent) bond primary bonds ~ ionic or covalent

polar bond ∆ in electronegativity of atoms EN ~ ability to attract e

EN value not absolutebut relative

Ch 1 #13

polar bond, polar molecule, dipole dipole moment µ = e d e ~ (partial) charge on the atom

d ~ distance betw charges

Read p12 and solve Prob 8 and 9!

Ch 1 #14

Atomic radius and electronegativity

Both determined by nuclear charge (# of protons)

# of shells (position of e)

Ch 1 #15

(electrostatic) potential map e distribution, (partial) charge distribution

size and shape of molecule

reactivity and reactive site (relative) size of Hin LiH, H2, HF

Ch 1 #16

Covalent ions covalent compounds containing charge ion no ionic bond, though

charge = # of protons - # electrons for NH4, 7 + 4 – 10 = 1

neutral neutral covalent ionmolecule molecule ion

Ch 1 #17

Formal charge charge assigned to an atom in a molecule charge distribution among atoms in a covalent species (ionic

or neutral)

FC = [# of valence e] – [# of e it owns]= [# val e] – [# non-bonding e + ½ (# bonding e)]

non-bonding electrons = lone-pair e’s = unshared pair of e’s

Ch 1 #18

Octet rule, formal charge, and stability CH4 vs CH3 radical

HCN vs HNC

H:C:::N: :C:::N:H

H

CH

H

H

C HH

H

satisfying octet rule NOT satisfying octet rulestable [not reactive] unstable [(very) reactive]

satisfying octet rule satisfying octet ruleformal charge ~ 0 0 0 formal charge ~ -1 +1 0stable unstable (actually, not likely present)

−:CH3

+CH3

Ch 1 #19

Valency # of bonds with no formal charge

if formal charge

for carbon carbocation ~ species containing C+

carbanion ~ species containing C-

radical ~ species w/ •

monovalent divalent trivalent tetravalent

oxonium ion

Ch 1 #20

Drawing Lewis structure Arrange the atoms. from center to peripheral ~ C, N, O, then X and H

Bond atoms (with e) satisfying octet rule.

Assign (formal) charge if needed.

Practice. Prob 14 p16

CH4O

C2H4

Ch 1 #21

Representing [drawing] structures Lewis structure ~ valence e’s

Kekule structure ~ bond as line, no : line(-bond) structure

condensed structure ~ no bond, if not necessary

skeletal structure ~ bonds only bond-line structure

a line for a bond; not showing C and H bonded to C

Section 2.6 p78

OH

Ch 1 #22

Table 1.5 p18

N

O

O

CH3CH2C(O)CH3

COOH

OH

O-C(=O)OH, -C(O)OH

-C(=O)H, -C(O)H

Ch 1 #23

Atomic orbitals (AO) AO describe the location of e (probability) density in atom quantum mechanics (Schrodinger eqn)

quantum numbers orbital

s orbitals

node where wave function is zero

no electron density

An e behaves like a standing wave.

+-

Ch 1 #24

p orbitals

3 p orbitals

lobe knob

Ch 1 #25

Molecular orbitals (MO) MO describes the location of e density in molecule combination of AO’s bonding MO

bond in particle sense

energy released= bond strength= bond dissociation energy

bond length

H2

Ch 1 #26

bond in wave sense conservation of orbitals ~ 2 AO’s 2 MO’s AO’s in-phase ~ reinforcing ~ overlap bonding MO

AO’s out-of-phase ~ cancelling ~ node antibonding MO

same e configuration in AO and in MO aufbau, exclusion

2 e in σ BMO, no e in σ* AMO σ bond

σ bond head-on overlap

H2

σ

σ∗

Ch 1 #27

bond order # of bonds betw atoms

(# bonding e - # antibonding e)/2

for H2, bond order = (2 – 0)/2 = 1 ~ single bond

Prob 20 p23 He2+ exist?

H2

σ

σ∗

Ch 1 #28

π bond side-to-side overlap

π BMO and π* AMO

weaker than σ bond

Prob 21 p25 σ, σ*, π, or π*?

π*

π

Ch 1 #29

Single bond and sp3 hybridization Experimental data for methane [CH4] shows 4 identical bonds

with tetrahedral geometry. tetrahedral VSEPR theory

VSEPR theory determines molecular shape.

≈ sawhorse drawing

Ch 1 #30

valence-shell e pair repulsion (VSEPR) theoryRule 1: VSEPs (bonding or non-bonding) repel each other

Rule 2: Lone pair repels more.

Rule 3: Double and triple bonds as one EP

Ch 1 #31

hybridization [混成化] ~ to bond well (or to explain well)

bonding 4 sp3-1s bonds repulsion tetrahedral

energy required

energy released

energy out > in bond thru hybridization

club

VSEPR-1

Ch 1 #32

ethane [CH3CH3]

bond angles all 109.5º?not exactly (111.5º and 107.5º)

Ch 1 #33

Double bond and sp2 hybridization Ethene [ethylene, CH2=CH2] is planar.

hybridization

bonding

VSEPR-3

Ch 1 #34

double bond 1 σ + 1 π

shorter and stronger than single bond C=C ~ 1.33 Å, 174 kcal/mol

C-C ~ 1.54 Å, 90 kcal/mol

174 < 90 x 2 π ~ side-to-side overlap ~ less overlap than σ

π bond still strong enough to prevent = from rotating

box p31 allotropes of carbon

diamond, graphite, CNT, fullerene, graphene

sp3 vs sp2

174 = 90 + 84?

Ch 1 #35

Triple bond and sp hybridization Ethyne [acetylene, CH≡CH] is linear.

hybridization

bonding triple bond

1.20 Å, 231 kcal/mol

Ch 1 #36

Carbon with 3 bonds methyl cation [+CH3] 3 sp2-s σ bonds + 1 empty p orbital ~ VESPR-1

methyl radical [•CH3] 3 sp2-s σ bonds + 1 p orbital w/ 1 e ~ VESPR-1

Ch 1 #37

methyl anion [−:CH3] 3 sp3-s σ bonds + 1 lone pair ~ VESPR-1

Ch 1 #38

H2O bond angle close to tetrahedral angle, not 180º

hybridization

104.5º < 109.5º VSEPR-2: non-bonding EP repels more.

more diffuse [larger] e distribution than bonding EP

Ch 1 #39

NH3 and NH4+

NH3

similar to H2O

104.5 < 107.3 < 109.5 1 lone pair EP ~ VSEPR-2

NH4+

4 identical bonds ~ 4 sp3

Ch 1 #40

Hydrogen halides [HX] X ~ halogen [F, Cl, Br, I]

sp3-s bond

bond length and strength the shorter, the stronger

H-Fmore overlap

H-Clless overlap

Ch 1 #41

Dipole moment of molecule is vector sum of bond (and lone-pair) dipoles

nonpolar

polar

‘Lone-pair dipole’ contributes.

NF3 µ = 0.24 D

Ch 1 #42

halomethanes

bond angles ~ explainable

dipole moments ~ rather complex

Cl

HH

H

Cl

HH

Cl

Cl

HCl

Cl108

µ = 1.54 D µ = 1.02 D

111

112112

110

Ch 1 #43

Summary: structure and bonding the shorter, the stronger the greater the e density in overlap

the more s character

(112+62)

(50% s)

(33% s)

(25% s)

Ch 1 #44

ACIDS and BASES Brφnsted-Lowry definition of acid and base acid ~ proton (H+) donor ~ HA HX, H2O, ROH, RNH2, RCH3

base ~ proton acceptor ~ B: RNH2, ROH, H2O (amphoteric), X-

acid-base rxn = proton transfer rxn ~ an equilibrium rxn

Equili moves to weaker acid (and base).

acid base base acid

Ch 1 #45

Acid strength ~ pKa

for acid HA in water

in dilute solution, [H2O] is constant at 55.5 M Ka ~ acid dissociation constant, acidity constant,

degree of ionization of acid in water

pKa ~ acidity strong acid ~ Ka > 1, pKa < 0

pH (of solution) vs pKa (of acid)

Ka = Keq [H2O] = [H3O+][A-]/[HA] or [H+][A-]/[HA]

pKa = – log Ka

strong and weak acid p43 ~ arbitrary

pKa of water?Prob 102 p67

Ch 1 #46

conjugate base, acidity, and equilibrium

Strong reacts to form weak.

The stronger the acid, the weaker the conjugate base. The more stable the conjugate base, the more reactive the acid.

stability vs reactivity

Ch 1 #47

Organic acids and bases common organic acid ~ carboxylic acid [RCOOH]

why? inductive + resonance effect

common organic base ~ amine [RNH2]

why? inductive effect

compared with ROH

NH4+

pKa = 9.25

CH3CH2NH2pKa > 40

pKb = 4.75 pKb = 3.3

Ch 1 #48

Strong and weak acids

HXinorgacids–10

–1.74

–615.74

RNH2amines35-40

RHhydrocarbons25-50

–2

5

strong acid(≈100%)

weak acid(partially ionized in water)

very weak acid (≈0%)

ROϴ, HOϴ

alkoxide ionhydroxide ion

RCOOϴ

carboxy anionXϴ

halide ion

Rϴ

RNHϴ

R ~ alkylCH3, CH3CH2, --

weak base medium base strong base

OH

10

Ch 1 #49

*Curved arrow ~ electron movement Curved arrows show e movement

reaction mechanism

thermodynamics kinetics mechanism

Ch 1 #50

Acidity: effect of atom bonded to H The weaker [more stable] the conjugate base,

the stronger the acid.

effect of EN

CH3OH vs CH3NH2?

More EN atom accommodates (-) charge better.

Ch 1 #51

effect of size

CH3OH vs CH3SH?

H loosely bound to larger atom. Larger atom accommodates (-) charge better. Size effect overweighs EN effect.

Ch 1 #52

effect of hybridization

Ch 1 #53

Acidity: inductive effect (of substituent)

EN X pulls e through σ bond better than H does

inductive e-withdrawing makes conj base stable acid strong

weaken O-H bond better leaving H acid strong

H < Br < Cl < F

e-withdrawing groups ~ most ~ -CN, -X, -OR, -C=O, etc

substitution reaction ~ 置換反應 substituent ~ 치환기

Ch 1 #54

R [alkyl] pushes e better than H does

inductive e-donating makes conj base strong acid weak

e-donating groups ~ -R, -O-, -COO-, etc

pKa ~ 13.1 pKa ~ 4.9 inductive effect (only)?

Ch 1 #55

Acidity: resonance effect (of subs) carboxylic acid vs alcohol

RCOOH is much stronger acid than ROH inductive effect of C=O

resonance effect ~ stabilize conj base by e delocalization

Ch 1 #56

pH and pKa

Henderson-Hasselbalch eqn

from definition of pH and pKa

tells [acidic form, HA]/[basic form, A-]in solution of certain pH acidic form when pH < pKa

basic form when pH > pKa

Ch 1 #57

useful for separation

at pH = 2, both in acidic form RCOOH to ether, RNH3

+ to water

charged (ionic) to water (polar);neutral (organic) to ether (organic).

|pH – pKa| > 2 for better separation (< 1/100)

Q. What pH for amine in ether and acid in water?

Do Prob 103 and 104 p67

pKa = 5

pKa = 10

Ch 1 #58

Buffer solution buffer solution with weak acid and its conj base

when ϴOH or H+ added no change in pH

preparation using H-H eqn

for pH, pKa, and buffer, Study guide pp36-55

RCOOH RCOOϴ + H+

RCOONa RCOOϴ + Na+

Ch 1 #59

Lewis acids and bases Lewis acid ~ accepts (a share in) an electron pair

Lewis base ~ donates (a share in) an electron pair

Lewis bases are BL bases.

Lewis acids are not limited to BL acids. Protonic acids are Lewis acids. ~ H+ accept e pair.

AlBr3, BF3, FeCl3, BH3, etc ~ usually-called Lewis acid do not give out H+; accepts e pair with empty orbital

Ch 1 #60

Exercise Do (all the) problems!

Prob 95. Which N more basic?

N

pyrrolidinebasic

pyrrolemuch less basic

N NN

NH H H

imidazolebasic

pyridinebasic

sp2sp3

N

Ch 1 #61