CHE-310
Organic Chemistry I
Dr. James Lyle; office: NSM D-323
(310) 243-3388 or 243-3376
jlyle@csudh.edu
office hours: MWF: 9-10:00am & Tu: 8:00-9:00am
Web page:
http://chemistry.csudh.edu/
texts:
Organic Chemistry,
Morrison & Boyd (6th)
Supplement to...,
Morrison & Boyd
(optional)
Supplement...
Model kit
Grading: traditional, no curve!
A=100%-93%, A-=92%-90%,B+=89%-88%,
B=87%-83%,etc.
Daily exams
4 exams @ 100 pts
final exam
homework
=
100
=
400
=
100
required
600
Daily exams
No make ups! Drop two lowest scores. Begin at 10:00!
Daily Homework: Required!
(hold until called for)
Cheating: Don’t do it! The penalties are severe.
Turn off all cell phones and pagers!
Organic Chemistry; difficult, challenging!
“memorization course” (NOT! well…maybe),
body of knowledge + application of theory!
How to succeed?
1. look over the text before lecture.
2. listen carefully to lectures
3. read the text (take notes)
4. do the homework (twice...?)
5. review
Organic Chemistry - the study of the compounds of carbon,
their properties and the changes that they undergo.
Descriptive approach nomenclature
syntheses
reactions
mechanisms
...
First: review topics from gen. chem. important to o-chem.
atomic structure
subatomic particles:
mass
charge
protons
1 amu
+1
neutrons
1 amu
0
electrons
~0 amu
-1
nucleus: protons & neutrons
electron shells & subshells:
electrons
atomic number = number of protons in the nucleus of the
atom (different for each element); Hydrogen = 1, Helium =
2, Lithium = 3,...
[also the number of electrons in a neutral atom]
Iron = 26
26 protons = +26
26 electrons=-26
net charge=
0
atomic mass = mass of an atom; sum of the weights of
the protons & neutrons.
But, not all atoms of a given element are identical.
isotopes - atoms of the same element with different
numbers of neutrons.
examples of isotopes
prot.
neut.
%
H1
H2
1
1
0
1
99.985
0.015
C12
C13
6
6
6
7
98.89
1.11
C14
6
8
...
Cl35
17
18
75.53
Cl37
17
20
24.47
F19
9
10
100
atomic weight: weighted average mass of the atoms;
combining weight...
electrons => energy shells & subshells about the nucleus.
shells = 1, 2, 3, 4, ...
subshells = s, p, d, f
orbitals = region in space where an electron of given energy
is likely to be found; no more than two electrons of opposite
spin per orbital (Pauli exclusion principle).
maximum number of electrons per
subshell:
s
2
p
6
d
10
f
14
order of filling
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
6f
spectral notation: 1s2,2s2,2p6…
Fluorine (at.# 9) 9p/9e
1s2,2s2,2p5
Chlorine (at.# 17) 17p/17e
1s2,2s2,2p6,3s2,3p5
Bromine (at.# 35) 35p/35e
1s2,2s2,2p6,3s2,3p6,4s2,3d10,4p5
Iodine (at.# 53) 53p/53e
1s2,2s2,2p6,3s2,3p6,4s2,3d10,4p6,5s2,4d10,5p5
valence electrons = electrons in the outermost shell
Fluorine has 7 valence elect.
Chlorine has 7 valence elect.
Bromine has 7 valence elect.
Iodine has 7 valence elect.
PERIODIC CHART OF THE ELEMENTS
I
VIII
┌────┐
┌────┐
│ H │
│ He │
│ 1 │ II
III
IV
V
VI VII │ 2 │
├────┼────┐
┌────┬────┬────┬────┬────┼────┤
│ Li │ Be │
│ B │ C │ N │ O │ F │ Ne │
│ 3 │ 4 │
│ 5 │ 6 │ 7 │ 8 │ 9 │ 10 │
├────┼────┤
├────┼────┼────┼────┼────┼────┤
│ Na │ Mg │
│ Al │ Si │ P │ S │ Cl │ Ar │
│ 11 │ 12 │
│ 13 │ 14 │ 15 │ 16 │ 17 │ 18 │
├────┼────┼────┬────┬────┬────┬────┬────┬────┬────┬────┬────┼────┼────┼────┼────┼────┼────┤
│ K │ Ca │ Sc │ Ti │ V │ Cr │ Mn │ Fe │ Co │ Ni │ Cu │ Zn │ Ga │ Ge │ As │ Se │ Br │ Kr │
│ 19 │ 20 │ 21 │ 22 │ 23 │ 24 │ 25 │ 26 │ 27 │ 28 │ 29 │ 30 │ 31 │ 32 │ 33 │ 34 │ 35 │ 36 │
├────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┤
│ Rb │ Sr │ Y │ Zr │ Nb │ Mo │ Tc │ Ru │ Rh │ Pd │ Ag │ Cd │ In │ Sn │ Sb │ Te │ I │ Xe │
│ 37 │ 38 │ 39 │ 40 │ 41 │ 42 │ 43 │ 44 │ 45 │ 46 │ 47 │ 48 │ 49 │ 50 │ 51 │ 52 │ 53 │ 54 │
├────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┤
│ Cs │ Ba │ La │ Hf │ Ta │ W │ Re │ Os │ Ir │ Pt │ Au │ Hg │ Tl │ Pb │ Bi │ Po │ At │ Rn │
│ 55 │ 56 │ 57 │ 72 │ 73 │ 74 │ 75 │ 76 │ 77 │ 78 │ 79 │ 80 │ 81 │ 82 │ 83 │ 84 │ 85 │ 86 │
├────┼────┼────┼────┼────┼────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┘
│ Fr │ Ra │ Ac │
│
│
│ 87 │ 88 │ 89 │104 │105 │
└────┴────┴────┴────┴────┘
┌────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┐
│ Ce │ Pr │ Nd │ Pm │ Sm │ Eu │ Gd │ Tb │ Dy │ Ho │ Er │ Tm │ Yb │ Lu │
│ 58 │ 59 │ 60 │ 61 │ 62 │ 63 │ 64 │ 65 │ 66 │ 67 │ 68 │ 69 │ 70 │ 71 │
├────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┤
│ Th │ Pa │ U │ Np │ Pu │ Am │ Cm │ Bk │ Cf │ Es │ Fm │ Md │ No │ Lr │
│ 90 │ 91 │ 92 │ 93 │ 94 │ 95 │ 96 │ 97 │ 98 │ 99 │100 │101 │102 │103 │
└────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┘
periodic chart of the elements
metals & nonmetals
families (groups) of elements
alkali metals (group I)
Li,Na,K,...
alkaline earths (group II)
Be,Mg,Ca,...
halogens (group VII)
F,Cl,Br,I,...
noble gases (group VIII or 0)
He,Ne,Ar,...
group number = valence elec.
Chemical bonding (classical)
chemical bond: force that holds atoms together in compounds.
ionic bond ~ between
metals & non-metals
covalent bond ~ between
non-metals & non-metals
definitions:
ionic bond: a chemical bond formed by the transfer of
valence electrons to achieve noble gas electron configurations, resulting in ions held together by electrostatic
attraction.
covalent bond: chemical bond formed by the sharing of
valence electrons to achieve noble gas electron
configurations.
ionic bond example:
sodium chloride
sodium = Na, atomic # 11
1s2,2s2,2p6,3s1
neon = Ne, atomic # 10
1s2,2s2,2p6
if Na loses 1 elect. then it will have a noble gas elect.
config. like Ne but will be charged, +1 ( 11p/10e ).
=> Na+
sodium ion
chlorine = Cl, atomic # 17
1s2,2s2,2p6,3s2,3p5
argon = Ar, atomic # 18
1s2,2s2,2p6,3s2,3p6
if chlorine can gain an electron it will have a noble gas
electron config. like argon but will be charged -1
(17p/18e) Clsodium chloride = NaCl
or
Na+Cl-
covalent bonds
Lewis Dot representations
.
. Be .
..
:..Cl.
Ne
.
.C .
.
..
. O...
H
H2O
=
..
H:O
..:H
see homework! review your gen chem text!
CO2
..
..
:O::C::O:
..
..
:O=C=O:
N2
:N:::N:
HCN
H:C:::N:
H-CN:
H2CO
..
H:C::O:
..
H
..
H-C=O:
|
H
:NN:
atomic orbitals
s
p
d
etc.
hybrid atomic orbitals
s
+
p
+
s + p + p => 3 sp2
s + p + p + p => 4 sp3
=>
2
sp hybrids
+
Hybrid atomic orbitals:
sp = linear; 180o
B A B
sp2 = trigonal; 120o
B A B
B
sp3 = tetrahedral; 109.5o
B
A B
B
B
VSEPR (valence shell electron pair repulsion)
prediction of hybridization
number of ligands (X)
plus
number of unshared pair of valence electrons (E)
equals
number of orbitals needed
 what type of hybrid orbitals are needed
eg. H2O =>
..
H:..
O:H
or
..
H—O—H
..
2 ligands + 2 lone pair = 4 orbitals
AX2E2
sp3 tetrahedral, 109.5o
H
H
O
water is a bent molecule with bond angles of 105o
VSEPR
AX2
sp
180o
linear
AX3
sp2
120o
trigonal
AX2E
sp2
~120o
or “bent”
AX4
sp3
109.5o
tetrahedral
AX3E
sp3
~109.5o
or “pyramidal
AX2E2
sp3
~109.5o
or “bent”
We can use the VSEPR method to predict the shape and
bond angles for simple covalent molecules.
SHAPE is important!
review gen chem text!
Do the homework!!!!!
PERIODIC CHART OF THE ELEMENTS
I
VIII
┌────┐
┌────┐
│ H │
│ He │
│ 1 │ II
III
IV
V
VI VII │ 2 │
├────┼────┐
┌────┬────┬────┬────┬────┼────┤
│ Li │ Be │
│ B │ C │ N │ O │ F │ Ne │
│ 3 │ 4 │
│ 5 │ 6 │ 7 │ 8 │ 9 │ 10 │
├────┼────┤
├────┼────┼────┼────┼────┼────┤
│ Na │ Mg │
│ Al │ Si │ P │ S │ Cl │ Ar │
│ 11 │ 12 │
│ 13 │ 14 │ 15 │ 16 │ 17 │ 18 │
├────┼────┼────┬────┬────┬────┬────┬────┬────┬────┬────┬────┼────┼────┼────┼────┼────┼────┤
│ K │ Ca │ Sc │ Ti │ V │ Cr │ Mn │ Fe │ Co │ Ni │ Cu │ Zn │ Ga │ Ge │ As │ Se │ Br │ Kr │
│ 19 │ 20 │ 21 │ 22 │ 23 │ 24 │ 25 │ 26 │ 27 │ 28 │ 29 │ 30 │ 31 │ 32 │ 33 │ 34 │ 35 │ 36 │
├────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┤
│ Rb │ Sr │ Y │ Zr │ Nb │ Mo │ Tc │ Ru │ Rh │ Pd │ Ag │ Cd │ In │ Sn │ Sb │ Te │ I │ Xe │
│ 37 │ 38 │ 39 │ 40 │ 41 │ 42 │ 43 │ 44 │ 45 │ 46 │ 47 │ 48 │ 49 │ 50 │ 51 │ 52 │ 53 │ 54 │
├────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┤
│ Cs │ Ba │ La │ Hf │ Ta │ W │ Re │ Os │ Ir │ Pt │ Au │ Hg │ Tl │ Pb │ Bi │ Po │ At │ Rn │
│ 55 │ 56 │ 57 │ 72 │ 73 │ 74 │ 75 │ 76 │ 77 │ 78 │ 79 │ 80 │ 81 │ 82 │ 83 │ 84 │ 85 │ 86 │
├────┼────┼────┼────┼────┼────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┘
│ Fr │ Ra │ Ac │
│
│
│ 87 │ 88 │ 89 │104 │105 │
└────┴────┴────┴────┴────┘
┌────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┐
│ Ce │ Pr │ Nd │ Pm │ Sm │ Eu │ Gd │ Tb │ Dy │ Ho │ Er │ Tm │ Yb │ Lu │
│ 58 │ 59 │ 60 │ 61 │ 62 │ 63 │ 64 │ 65 │ 66 │ 67 │ 68 │ 69 │ 70 │ 71 │
├────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┤
│ Th │ Pa │ U │ Np │ Pu │ Am │ Cm │ Bk │ Cf │ Es │ Fm │ Md │ No │ Lr │
│ 90 │ 91 │ 92 │ 93 │ 94 │ 95 │ 96 │ 97 │ 98 │ 99 │100 │101 │102 │103 │
└────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┘
Polarity
Covalent bonds are polar when the two atoms sharing
electrons have different electronegativities.
eg.
H—Cl
δ+ δ-
a charge separation or a dipole gives a polar bond.
O2
.. ..
:O=O: has a non-polar bond
Representation of dipoles using vectors
a) magnitude = length
b) direction = positive  negative
A molecule will be non-polar if the vector sum of the bond
dipoles is zero; eg. they cancel one another.
A molecule with be polar if the vector sum of the bond
dipoles is non-zero.
Determining polarity of covalent molecules:
1. Lewis dot structure
2. VSEPR  hybridization  shape of the molecule
3. dipoles for polar bonds
4. vector sum of the bond dipoles
5. vector sum = 0  non-polar molecule
6. vector sum  0  polar molecule
CO2
:O=C=O:
sp linear
vector sum = 0
non-polar molecule
H2O
..
H—O—H
..
H O
H
AX2E2
sp3
vector sum  0
polar molecule!
tetrahedral (bent)
CH3OH
Both C & O are sp3
hybridized.
The bond dipole vectors do
not cancel each other and
the molecule is polar.
NB: must know shape to
determine polarity!
H
C H
O
H
H
Intermolecular forces. Attractions between molecules.
ionic attractions
(very strong)
Na+ClCl-Na+
dipole-dipole attractions
H—Br
Br—H
hydrogen bonding ( H attached to N,O,F )
H—O----H—O
|
|
H
H
van der Waals (London forces)
(weak)
Br—Br
Br—Br
intermolecular attractions
strongest
ionic attractions
dipole-dipole / hydrogen bonding
van der Waals
weakest
ionic bonds => ionic attractions
polar covalent => dipole-dipole attractions
non-polar covalent => van der Waals
Cl2
non-polar covalent => van der Waals
CO2
non-polar covalent => van der Waals
H2O
polar covalent => dipole-dipole &
Hydrogen bonding
CH4
non-polar covalent => van der Waals
KBr
ionic bonding => ionic attractions
bonding => shape => polarity => physical properties
physical properties:
melting point
boiling point
solubility
The stronger the intermolecular forces the higher the mp/bp.
Ionic substances have significantly higher mp/bp than do
covalent substances. [note: mp/bp also increase with
increasing size.]
Prediction of mp/bp (relatively high or low?):
mp
bp
Mg(OH)2
ionic => ionic attractions
350oC --
CH3OH
polar => dipole-dipole + H-bond
-94oC 65oC
CH2O
polar => dipole-dipole
-920C -21oC
CH3CH3
non-polar => van der Waals
-183oC –89oC
Solubility
“like dissolves like”
~ water soluble? must be ionic or highly polar + H-bond
(hydrophilic)
~ water insoluble? must be non-polar or weakly polar
(hydrophobic)
Most organic compounds are water insoluble!
Acids/Bases
historic:
acids – from L. acidus = “sour”
sour taste
react with metals  H2
react with bases  water + salts
change litmus  red
react with limestone  CO2
examples: HCl, H2SO4, HNO3, HClO4
historic:
bases - bitter taste
soapy feel
react with acids  water + salts
change litmus  blue
examples: NaOH, Al(OH)3, K2CO3, NaHCO3
Lowry-Brønsted Acid - a substance that donates a proton
(H+) in a chemical reaction.
Lowry-Brønsted Base – a substance that accepts a proton
(H+) in a chemical reaction.
CH3MgBr + NH3  CH4 + Mg(NH2)Br
base
NaOH
base
acid
+ H2SO4
acid
acid

H2O
acid
base
+ NaHSO4
base
Lewis Acid – a substance that accepts an electron pair in a
chemical reaction to form a covalent bond.
Lewis Base – a substance that donates an electron pair in a
chemical reaction to form a covalent bond.
BF3 + :NH3
Lewis
- +
 F3B:NH3
Lowry-Brønsted
Rule: acid/base reactions must run “down hill.”
stronger acid/base  weaker acid/base
H2SO4 + H2O 
stronger
stronger
acid
base
H2O +
weaker
acid
NH3 
weaker
base
HSO4- +
weaker
base
NH4+ +
stronger
acid
(note direction of reactions)
H3O+
weaker
acid
OHstronger
base
Within a period of the periodic chart, acid strength increases with
increasing electronegativity:
CH4 < NH3 < H2O < HF
Within a family of elements, acid strength increases with
increasing size:
HF < HCl < HBr < HI
PERIODIC CHART OF THE ELEMENTS
I
VIII
┌────┐
┌────┐
│ H │
│ He │
│ 1 │ II
III
IV
V
VI VII │ 2 │
├────┼────┐
┌────┬────┬────┬────┬────┼────┤
│ Li │ Be │
│ B │ C │ N │ O │ F │ Ne │
│ 3 │ 4 │
│ 5 │ 6 │ 7 │ 8 │ 9 │ 10 │
├────┼────┤
├────┼────┼────┼────┼────┼────┤
│ Na │ Mg │
│ Al │ Si │ P │ S │ Cl │ Ar │
│ 11 │ 12 │
│ 13 │ 14 │ 15 │ 16 │ 17 │ 18 │
├────┼────┼────┬────┬────┬────┬────┬────┬────┬────┬────┬────┼────┼────┼────┼────┼────┼────┤
│ K │ Ca │ Sc │ Ti │ V │ Cr │ Mn │ Fe │ Co │ Ni │ Cu │ Zn │ Ga │ Ge │ As │ Se │ Br │ Kr │
│ 19 │ 20 │ 21 │ 22 │ 23 │ 24 │ 25 │ 26 │ 27 │ 28 │ 29 │ 30 │ 31 │ 32 │ 33 │ 34 │ 35 │ 36 │
├────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┤
│ Rb │ Sr │ Y │ Zr │ Nb │ Mo │ Tc │ Ru │ Rh │ Pd │ Ag │ Cd │ In │ Sn │ Sb │ Te │ I │ Xe │
│ 37 │ 38 │ 39 │ 40 │ 41 │ 42 │ 43 │ 44 │ 45 │ 46 │ 47 │ 48 │ 49 │ 50 │ 51 │ 52 │ 53 │ 54 │
├────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┤
│ Cs │ Ba │ La │ Hf │ Ta │ W │ Re │ Os │ Ir │ Pt │ Au │ Hg │ Tl │ Pb │ Bi │ Po │ At │ Rn │
│ 55 │ 56 │ 57 │ 72 │ 73 │ 74 │ 75 │ 76 │ 77 │ 78 │ 79 │ 80 │ 81 │ 82 │ 83 │ 84 │ 85 │ 86 │
├────┼────┼────┼────┼────┼────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┘
│ Fr │ Ra │ Ac │
│
│
│ 87 │ 88 │ 89 │104 │105 │
└────┴────┴────┴────┴────┘
┌────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┐
│ Ce │ Pr │ Nd │ Pm │ Sm │ Eu │ Gd │ Tb │ Dy │ Ho │ Er │ Tm │ Yb │ Lu │
│ 58 │ 59 │ 60 │ 61 │ 62 │ 63 │ 64 │ 65 │ 66 │ 67 │ 68 │ 69 │ 70 │ 71 │
├────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┤
│ Th │ Pa │ U │ Np │ Pu │ Am │ Cm │ Bk │ Cf │ Es │ Fm │ Md │ No │ Lr │
│ 90 │ 91 │ 92 │ 93 │ 94 │ 95 │ 96 │ 97 │ 98 │ 99 │100 │101 │102 │103 │
└────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┘
Which is the stronger acid?
H2O or H2S?
oxygen & sulfur are in the same family and sulfur is
bigger: H2S > H2O
What is the order of base strength?
F-
Cl-
Br-
I-
in the halogen family base strength decreases with
increasing size:
F- >
Cl- > Br- > I-
Will H2O react with NaSH as shown below?
H2O
+
NaSH

NaOH
+
WA
H2S
SA
no, H2O < H2S
Will the following reaction proceed as shown?
HI
+
NaCl
SA
yes, HI > HCl

HCl
WA
+
NaI
Isomers - different compounds with the same molecular
formula.
example: C2H6O
CH3CH2OH
CH3OCH3
ethyl alcohol
dimethyl ether
bp 78oC
bp –24oC
H H
H C C O H
H H
H
H
H C O C H
H
H