PART 1
Introduction to
Organic Chemistry
Chapter
1. Structure, bonding, acid-base
2. Nomenclature, physical properties,
drawing structure
Chapter 1
Structure and Bonding Acids and Bases
Electronic structure Chemical bonds
Acid-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, CO
2, Na
2CO
3, 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 bonding acid-base
2. nomenclature, physical properties structure 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 atom
Define 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 e
to 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 absolute but 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 H in 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 ion molecule 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
CH
4vs CH
3radical
HCN vs HNC
H:C:::N: :C:::N:H
H H C
H H
C H H
H
satisfying octet rule NOT satisfying octet rule stable [not reactive] unstable [(very) reactive]
satisfying octet rule satisfying octet rule
formal charge ~ 0 0 0 formal charge ~ -1 +1 0
stable unstable (actually, not likely present)
−
:CH
3+
CH
3Ch 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
-C(=O)OH, -C(O)OH O -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 sp 3 hybridization
Experimental data for methane [CH
4] 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) theory
Rule 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 [CH
3CH
3]
bond angles all 109.5º?
not exactly
(111.5º and 107.5º)
Ch 1 #33
Double bond and sp 2 hybridization
Ethene [ethylene, CH
2=CH
2] 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 [
+CH
3]
3 sp2-s σ bonds + 1 empty p orbital ~ VESPR-1
methyl radical [
•CH
3]
3 sp2-s σ bonds + 1 p orbital w/ 1 e ~ VESPR-1
Ch 1 #37
methyl anion [
−:CH
3]
3 sp3-s σ bonds + 1 lone pair ~ VESPR-1
Ch 1 #38
H 2 O
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
NH 3 and NH 4 +
NH
3 similar to H2O
104.5 < 107.3 < 109.5 1 lone pair EP ~ VSEPR-2
NH
4+ 4 identical bonds ~ 4 sp3
Ch 1 #40
Hydrogen halides [HX]
X ~ halogen [F, Cl, Br, I]
sp
3-s bond
bond length and strength
the shorter, the stronger
H-Fmore overlap H-Cl
less 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
H H
H
Cl
H H
Cl
Cl
H Cl
Cl
108
µ = 1.54 D µ = 1.02 D
111
112 112
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 ~ p K a
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 [RNH
2]
why? inductive effect
compared with ROH
NH4+ pKa = 9.25
CH3CH2NH2 pKa > 40 pKb = 4.75 pKb = 3.3
Ch 1 #48
Strong and weak acids
HX
inorg acids –10
–1.74
–6 15.74
RNH2
amines 35-40
RH
hydrocarbons 25-50
–2
5
strong acid (≈100%)
weak acid
(partially ionized in water)
very weak acid (≈0%)
ROϴ, HOϴ
alkoxide ion hydroxide ion
RCOOϴ
carboxy anion
Xϴ
halide ion
Rϴ RNHϴ R ~ alkyl
CH3, 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 p K a
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, p K
a, 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
pyrrolidine
basic pyrrole
much less basic
N N
N
H H N H
imidazole
basic pyridine
basic sp2
sp3
N
Ch 1 #61
