Organic Chemistry:
The study of carbon and its
compounds
Partial Periodic Table
Know these elements!
Group IA
Period
IIA
IIIA
IVA
VA
VIA
VIIA
N
O
F
P
S
Cl
H
1
2
Li
Be
B
3
Na
Mg
Al
4
K
Ca
See the inside back cover of text
Si
Br
Organization of Chemical
Compounds
Compound
Organic
Contains C & H,
may have N, O, S
Covalent bonds
Organometallic
C covalently
bonded to
metal, i.e. Mg
Inorganic
All elements
possible
Ionic Bonds
1. Bonding:
Ionic Bonding—transfer of electrons.
So:
1. Bonding
(continued):
Covalent Bonding—sharing of
electrons.
H shares
H + H
H ..H Each
two electrons.
The simplest way to symbolize the bonding
of a covalent molecule is to use Lewis dot
structures. In doing so, it is important to
remember the octet rule.
1. Bonding (continued):
Octet Rule:
Atoms will react (i.e. gain, lose or
share electrons) in order to have the
same number of valence electrons as
the nearest noble gas. The sum of all
the shared and unshared valence
electrons about an atom must total 8.
Note that this does not apply to all
atoms (e.g. H, B, Al and sometimes S.
1. Bonding (continued)
Lewis/Kekule Structures
dots = electrons
line = two shared electrons
all unshared electrons are shown, and
are referred to as nonbonding
electrons. Each pair of non-bonding
electrons is also known as a “lone pair”
H
H
CH3Cl
H C Cl
H
Lewis
H
C
H
Kekule
Cl
1. Bonding (continued):
Note: If you don’t remember how to determine whether
or not a Lewis structure has a formal charge, you should
review it immediately!!!
Multiple bonding in Lewis/Kekule
structures:
Single Bond: sharing of one electron
pair between two atoms.
Double Bond: sharing of two electron
pairs.
Triple Bond: sharing of three electron
pairs.
1. Bonding (continued):
H
H
C
Lewis
H
H
C
C
C
H
H
H
H
Kekule
H
C
H
H
C
C
H
ethylene
O
H
H
C
C
H
C
O
H
acetylene
formaldehyde
Now, try to do Problems 1-2 and 1-3 on pages 8 and 9
of the text.
1. Bonding (continued):
Bonding Patterns
# of bonds
C
N
O
X
4
3
2
1
H
1
P
3 or 5
S
2, 4 or 6
Structural Formulas: There are two types of structural
formulas; complete Lewis structures and condensed
formulas.
See boxes pages 9 & 13
1. Bonding (continued):
Condensed Structural Formulas
You must be able to use these by next
week
Are written without showing all the
individual bonds
Each central atom is shown together
with the atoms bonded to it
Multiple bonds are drawn as they would
be in a Lewis structure
See pages 17-19
1. Bonding (continued):
Line-Angle Structural Formulas
You must be able to use these by next
week
Also known as a skeletal structure or a
stick figure
Bonds are represented by lines
Carbon atoms are assumed to be
present wherever two lines meet and at
the beginning or the end of a line
1. Bonding (continued):
Condensed and Line-Angle
Formulas
H H H H H H
H C C C C C C H
H H H H H H
Kekule
CH3CH2CH2CH2CH2CH3 = CH3(CH2)4CH3 =
1 2 3 4 5 6
1 2 3 4 5 6
condensed
condensed
line-angle
You should know how to draw condensed structures
from line-angle structures and vice-versa. Do problem
1-10 on page 19 of the text.
Review:
1. How many carbons does this compound have? How many
hydrogens are at each of the carbons?
2. How many carbons does each of the compounds below
have? How many hydrogens are at each of the carbons?
2. Electronegativity
Electronegativity refers to the relative
tendency of an atom to draw electrons to
itself in a chemical bond.
Electronegativity increases from left to right
and bottom to top across the periodic chart,
with F being the most electronegative
element.
In general, when considering
electronegativity, our reference point will be
carbon.
Electronegativities of the elements
Group IA
Period
1
IIA
IIIA
IVA
VA
VIA
VIIA Noble
Gases
H
2.1
H
2.1
He
2
Li
1.0
Be
1.6
B
1.8
C
2.5
N
3.0
O
3.4
F
4.0
Ne
3
Na
0.9
Mg
1.3
Al
1.6
Si
1.9
P
2.2
S
2.6
Cl
3.2
Ar
4
K
0.8
Ca
1.0
Ga
1.6
Ge
1.8
As
2.0
Se
2.4
Br
3.0
Kr
See page 10
2. Electronegativity (continued):
D Electronegativity good polarity
approximation
Know trend & approximate D values
C and H
these elements are
less electronegative
than carbon
these elements are
more electronegative
than carbon
periodic table
2. Electronegativity (continued):
Electronegativity and Bond
Polarity
A covalent bond with electrons shared
equally between the two atoms is called a
nonpolar bond.
A covalent bond in which the electrons are
shared unequally is called a polar bond or a
polar covalent bond.
H
Nonpolar
covalent bond
H
C
Cl
Polar covalent
bond
+
Na Cl
Ionic bond
2. Electronegativity (continued):
Polar Covalent Bonds
In the carbon—chlorine bond, the electrons are
more strongly attracted to Cl than to C because Cl
is more electronegative than C. This makes the Cl
more ‘negative’, and the C more ‘positive. This is
depicted below by the use of a d (delta) symbol.
We say that C has a partial positive (d+) charge,
and Cl has a partial negative (d-) charge.
The bond polarity is symbolized by
an arrow with its head at the negative
end and a cross at the positive end of
the polar bond.
Do problem 1-5 on page 11 of the
text.
d
H
d
Cl
Bonds are on a continuum
from Ionic to Covalent
Carbon almost never forms fully ionic bonds
3. Resonance
Recall from your knowledge of general chemistry
the concept of formal charges. The formal charge
= (group number) - (nonbonding electrons) 1/2(shared electrons).
Example: Compute the formal charge at carbon and nitrogen
in the species [H2CNH2]+.
Step 1: Draw the Lewis
H
Structure. In doing so,
you will see that there
C
are two possible
Lewis structures.
H
Step 2: Calculate formal
charge.
H
H
and
N
H
H
C
H
N
H
3. Resonance
In the first Lewis structure, the formal charge for C
is: FCc = 4 - 0 - 4 = 0.
The formal charge for N is: FCN = 5 - 0 - 4 = +1
H
H
C
H
N
H
H
+
N
C
H
H
H
In the second Lewis structure, the formal charge for
C is : FCc = 4 - 0 - 3 = +1.
The formal charge for N is: FCN = 5 - 2 - 3 = 0.
H
Do problem
1-6 on page
14.
H
C
H
H
N
H
C
H
H
N
H
3. Resonance
When two or more Lewis structures are possible,
differing only in the placement of electrons (not
atoms), the molecule will show characteristics of
both structures. These different structures are
called resonance structures or resonance forms.
They are not different compounds; they are just
different ways of drawing the same compound.
Individual resonance forms do not exist. The
actual structure of the molecule is a combination of
all of its resonance forms, and is said to be a
resonance hybrid. In the true structure, the
positive charge is spread out over both the C and N
atoms. It is said to be delocalized over the C and N
atoms.
3. Resonance
The interconversion of resonance forms is represented by
a double headed arrow.
H
H
+
C
H
N
H
H
C
H
H
resonance forms
+
N
=
H
H
d
H
H
C
d+
N
H
resonance hybrid
You can see in the above example that the charge is spread,
or delocalized over two atoms: C and N. This makes the ion
more stable that in would be if the entire charge were
localized on only one atom. Therefore, we call this a
resonance-stabilized cation. Resonance promotes
stabilization!
3. Resonance
Although separate resonance forms do not exist, we can
estimate their energies if they did exist. More stable resonance
forms are closer representations of the real molecule than less
stable ones. The more stable resonance form is called the
major contributor, and the less stable resonance form is
called the minor contributor. The true structure of the
compound resembles the major contributor more than it does
the minor contributor. Resonance forms can be compared by
the following criteria, beginning with the most important:
1.
2.
3.
4.
As many octets as possible.
As many bonds as possible.
Any negative charges on electronegative atoms.
As little charge separation as possible.
Based upon these rules, let us determine which is the major
resonance contributor in the example that we just did.
3. Resonance
Note the curved arrow
notation which is used
in organic chemistry
to show the direction
and movement of
electrons.
Number of octets:
Number of bonds:
Charge separation?:
- Charge on E atoms?:
B
A
H
H
+
C
H
N
H
H
C
H
+
N
H
1
5
No
No
H
2
6
No
No
Thus, all things being equal, we can see that structure B is the
major resonance contributor, since it has the greater number
of octets, and the greater number of bonds.
Do problems 1-7 and 1-8 on page 17 of the text.
Acidity and Basicity
There are three definitions of Acids and Bases: the Arrhenius
definition, the Bronsted-Lowry definition, and the Lewis
definition.
Arrhenius: An acid is a substance that dissociates in
water to give H3O+, and a base is a substance that
dissociates in water to give OH-.
Bronsted-Lowry: An acid is a species that can donate
a proton, and a base is a species that can accept a
proton.
Lewis: An acid is an electron pair acceptor, and a
base is an electron pair donor.
Let us now consider some examples to see whether we
can identify which definitions apply in specific cases.
4. Acidity and Basicity
H2SO4
HCl
BF3
+
+
+
H2O
HSO4
NaOH
NaCl
NH3
+
H3N
+
H3O
H2O
BF3
We can see from these examples that the Arrhenius and
Bronsted-Lowry definitions of acids and bases fail in
some instances in allowing us to identify which species
is an acid and which is a base. The Lewis definition is the
only one that will allow us to make the correct predictions
all the time.
Do not forget that when an acid reacts it forms a conjugate
base, and when a base reacts if forms a conjugate acid!
4. Acidity and Basicity
The strength of both Bronsted-Lowry and Arrhenius acids
is expressed by the extent of their ionization in water. The
general reaction of an acid (HA) with water is as follows:
HA
+
H2O
Ka
H3O +
A
where Ka represents the acid-dissociation constant.
+
-
[H3O ][A ]
Ka =
[HA]
The magnitude of Ka represents the relative strength of the
acid. The stronger the acid, the more it dissociates, giving
a larger Ka.
4. Acidity and Basicity
Strong Acids: Are almost completely ionized in water; Have
Ka’s greater than 1.
Weak Acids: (most organic compounds); have Ka’s less than
10-4.
Ka’s are often expressed on a logarithmic scale: pKa = -log10Ka
For water:
H2O
and
+
H2O
OH
+
H3O
-
-14
[H3O ][OH ]
-16
1.00 x 10
Ka =
=
= 1.8 x 10
[H2O]
55.6 mol/L
+
The logarithm of 1.8 x 10-16 is -15.7, and the pKa of water
is 15.7
See page 24 of the text.
4. Acidity/Basicity
Table 1-5 on page 25 of the text lists some common organic
and inorganic acids and their conjugate bases. From this
table, you can see that:
1.
2.
3.
4.
Strong acids generally have pKa values of ~ 0.
Weak acids generally have pka values of > 4.
The weaker the acid, the larger the pKa.
The weaker the acid, the smaller the Ka.
Predicting the direction of Acid/Base Equilibria:
In order to predict the direction of acid/base equilibria, it is
important to remember the following:
1. Acid base reactions favor formation of the weaker acid and
the weaker base.
2. The stronger the acid, the weaker its conjugate base.
3. The weaker the acid, the stronger its conjugate base.
4. Acidity/Basicity
Examples:
HCl
+
H2O
strong
acid
CH3-OH
weak
acid
-
Cl
+
+
H3O
weak
base
+
H2O
-
CH3O
+
strong
base
Do problems 1-14, 1-15 and 1-16 on page 26 and 28 of the text..
+
H3O
4. Acidity/Basicity
Predicting Acidity
Generally speaking, the more stable a conjugate base is, the
more acidic the acid from which it came. There are three
factors that affect the stability of conjugate bases. These are
electronegativity, size, and resonance.
Electronegativity: The more electronegative
an element is, the more easily it bears
a negative charge, giving a more
stable conjugate base,
and a stronger acid.
4. Acidity/Basicity
Electronegativity
C
<
N
<
O
<
F
electronegativity increases
Stability
Acidity
-
CH3
<
H-CH3
<
-
NH2
<
H-NH2 <
-
OH
<
H-OH
<
-
F
H-F
acidity increases
Basicity
-
CH3
>
-
NH2
>
-
OH
basicity increases
See page 29 of the text.
>
-
F
4. Acidity/Basicity
Size: The negative charge of an anion is more stable if is
spread over a large region of space. Within a column of
the periodic table, acidity increases down the column, as
the size of the elements increases.
H-F
<
Acidity
Stability
H-Cl <
H-Br
<
H-I
<
I
acidity increases
-
F
<
-
Cl
<
Br
size increases
-
-
4. Acidity/Basicity
Resonance Stabilization: A conjugate base of an acid is
more likely to form if it can be stabilized by resonance. See
bottom of page 29 in the text.
CH3CH2
O
ethoxide ion from ethanol
(pKa = 16)
O
H3C
C
O
O
H3C
C
O
acetate ion from acetic acid (pKa = 4.74)
O
O
H3C
S
O
O
H3C
S
O
O
O
H3C
S
O
O
methanesulfonate ion from methanesulfonic acid (pKa = -1.2)
4. Acidity/Basicity
One Final Note:
Non proton Lewis acids are known as Electrophiles.
Lewis Bases are known as Nucleophiles.
Most of the acid/base chemistry that we will encounter in
this course will be described in terms of nucleophiles and
electrophiles. It is very very important that you be
thoroughly familiar with these two terms.
Common Symbols &
Abbreviations
,e &e
Electron
Radical
Cation
Pair of
Electrons
Radical
Anion
Cation
Bonds
Anion
More abbreviations
D = difference
d = partial (d+, d-)

= polar bond, in direction of
electron polarization