Principles of drug Synthesis
.
Some common structures
Positions of substituents
Positions of aromatic substituents
Names and abbreviations for carbon
chains
Various alcohol types
Alkene, alkyne, alcohol, ether
Amine , nitro
How to draw Lewis’ structure correctly
Aldehyde, ketone, acid, ester
amide
oxidation levels
Various oxidation levels of carbon atom
A review of bonding theory
the ionic bonding force arises from the electrostatic
attraction between ions of opposite charge
the covalent bonding force arises from sharing of
electron pairs between atoms.
The Lewis structure notation

essential qualitative information about properties of
chemical compounds.

chemical properties of the groups that make up organic
molecules are to a first approximation constant from
molecule to molecule
an atom and its valence electrons

The element symbol represents the core
(the nucleus and all the inner-shell electrons)
The core carries a number of positive charges
+ charge = the number of valence electrons

The electrons are shown explicitly
Ions are obtained by adding or removing electrons

The charge on an ion = core charge - number of
electrons shown explicitly
A covalent bond model is constructed by allowing
atoms to share pairs of electrons.
Ordinarily, a shared pair is designated by a line:
C-C
H-H
.

All valence electrons of all atoms in the structure must
be shown explicitly.

Those electrons not in shared covalent bonds are
indicated as dots

For ion containing > 2 atoms covalently bonded to each
other 
the total charge on the ion =
the total core charge - the total number of electrons
(shared and unshared)
Lewis structures



all unshared electrons around the atom and all electrons in
bonds leading to the atom must be counted.
The valence-shell occupancy must not exceed 2 for hydrogen
and must not exceed 8 for atoms of the first row of the periodic
table.
For elements of the second and later rows, the valence-shell
occupancy may exceed 8.
Formal charge
- a bookkeeping device for electrons
- a rough guide to the charge distribution within a molecule
Assign to each atom all of its unshared pair electrons and
half of all electrons in bonds leading to it
 its electron ownership.

The formal charge of each atom =
core charge – electron ownership

The electron ownership of H is 1, its core
charge is + 1  its formal charge = zero

The electron ownership of oxygen is 7, and
the core charge is +6 formal charge = -1
.
All nonzero formal charges must be shown explicitly in the
structure.
the procedure in writing Lewis structures
1.
2.
3.
4.
5.
Count the total nb of valence electrons contributed by the
electrically neutral atoms.
for ion  add one e to the total for
each negative charge; subtract one for each positive charge.
Write the core symbols for the atoms and fill in the number of
electrons determined in Step 1. The electrons should be added so
as to make the valence- shell occupancy of hydrogen 2 and the
valence-shell occupancy of other atoms not less than 8 wherever
possible.
Valence-shell occupancy must not exceed 2 for H and 8 for a firstrow atom; for a second-row atom it may be 10 or 12.
Maximize the number of bonds, and minimize the number of
unpaired electrons, always taking care not to violate Rule 3.
Find the formal charge on each atom.
Lewis structure of NO2
a class of structures, for which the properties are not those
expected from the Lewis structure.

The thermochemical properties of various types of bonds are in
most instances transferable with good accuracy from molecule to
molecule

the heat of hydrogenation of benzene is less exothermic by about
37 kcal mole-1 than one would have expected from Lewis structure
on the basis of the measured heat of hydrogenation of ethylene
a discrepancy of this magnitude requires a fundamental modification of the bonding
model.
.

1
2
another Lewis structure of benzene, 2, is
identical to 1 except for the placement of the
double bonds.


The superposition of two or more Lewis structures into a
composite picture is called resonance
the term resonance tends to convey the idea
of a changing back and forth with time
 incorrect idea
difficult to avoid the pitfall of thinking of the benzene molecule as a
structure with three conventional double bonds, of the ethylene type,
jumping rapidly back and forth from one location to another



The electrons in the molecule move in a field of force created by
the six carbon and six hydrogen nuclei arranged around a regular
hexagon
Each of the six sides of the hexagon is entirely
equivalent to each other side;
there is no reason why electrons should, even momentarily, seek
out three sides and make them different from the other three, as
the two alternative 1 and 2 seem to imply that they do.
A less misleading picture: the circle in the middle of the ring
implies a distribution of the six double bond electrons of the
same symmetry as the arrangement of nuclei.
1

2
We shall continue to use the notation 1 and 2, as it has certain
advantages for thinking about reactions.
The most important features of structures for which resonance is
needed:
 the molecule is more stable (of lower energy) than one would expect
from looking at one of the individual structures,
 the actual distribution of electrons in the molecule is different from
what one could expect on the basis of one of the structures.
resonance is often referred to as delocalization.
 electrons are free to move over a large area of the molecule
Many other structures are of the necessity for modifying the Lewis
structure language by the addition of the resonance concept
The carboxylic acids are much stronger acids than the alcohols due
largely to greater stability of the carboxylate ion (6) over the
alkoxide (7 )
.
The allylic system may be cation (8), anion (9),
and radical (10), are all more stable than their
saturated counterparts.

there is for each an alternative structure
The rules in using resonance notation
1.
All nuclei must be in the same location in every structure.
2. Structures with nuclei in different locations, for example 15 and 16, are
chemically distinct substances, and interconversions between them are
actual chemical changes, always designated by
3. Structures with fewer bonds or with greater separation of formal charge are
less stable than those with more bonds or less less charge separation.
Thus 11 and 12 are higer-energy, respectively, than 13 and 14.
1.
Molecular Geometry
Lewis structures provide a simple method of estimating
molecular shapes.
The geometry about any atom covalently bonded to two
or more other atoms is found by counting the number
of electron groups around the atom.
Each unshared pair counts as one group, and each
bond, whether single or multiple, counts as one
group.
The number of electron groups around an atom is
therefore equal to the sum of the number of electron
pairs on the atom and the number of other atoms
bonded to it.
 .
The geometry



linear if the number of electron goups is two
trigonal if the number is three,
tetrahedral if the number is four
.

The rule is based on the electron-pair repulsion
model  electron pairs repel each other, they will
try to stay as far apart as possible.
if the electron groups are all equivalent
 the shape will be exactly trigonal (120o bond
angles),
 or exactly tetrahedral (109.5o bond angles),
BH3 or CH3+ (trigonal),
CH4 or NH4+ (tetrahedral).
.
If the groups are not all equivalent, the angles will deviate from the
ideal values.
Thus in NH3 (four electron-groups, three in N-H bonds, one an
unshared pair), the unshared pair, being attracted only by the
nitrogen nucleus, will be closer to the nitrogen on the average
than will the bonding pairs, which are also attracted by a
hydrogen nucleus.
Therefore the repulsion between the unshared pair and a bonding
pair is greater than between two bonding pairs, and the bonding
pairs will be pushed closer to each other.
The H-N-H angle should therefore be less than 109.5o (It is found
experimentally to be 107o)
in H2O (four electron groups, two unshared pairs, and two O-H
bonds), the angle is 104.5o.