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Organic Chemistry Notes
(UPMC, Licence Chimie, LC 204)
Introduction
The Chemical Bond: a revision.
Quantum Shells (1, 2, 3, 4, 5) and Sub-Orbitals (s, p, d, f)
Electrons in atoms arrange themselves in the quantum shells in groups of 2, 8, 8, 18, etc. There are four types
of orbitals (also called sub-levels), s, p, d, and f, which bear different shapes, each one being able to host up to
two electrons.
1st Quantum Shell
contains 2 electrons
1s (x 1) 2 electrons
2nd Quantum Shell
contains 8 electrons
2s (x 1) 2 electrons
2p (x 3) 6 electrons
3rd Quantum Shell
contains 8 electrons
3s (x 1) 2 electrons
3p (x 3) 6 electrons
4th Quantum Shell
contains 18 electrons
4s (x 1) 2 electrons
3d (x 5) 10 electrons
4p (x 3) 6 electrons
5th Quantum Shell
contains 18 electrons
5s (x 1) 2 electrons
4d (x 5) 10 electrons
5p (x 3) 6 electrons
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Table 1. Ground state electron configuration in elements
Atomic Number
Symbol
1
2
3
4
5
6
7
8
9
10
11
H
He
Li
Be
B
C
N
O
F
Ne
Na
Configuration
1s1
1s2
1s2 2s1
1s2 2s2
1s2 2s2 2p1
1s2 2s2 2p2
1s2 2s2 2p3
1s2 2s2 2p4
1s2 2s2 2p5
1s2 2s2 2p6
[Ne] 3s1
15
16
17
18
P
S
Cl
Ar
[Ne]
[Ne]
[Ne]
[Ne]
3s2
3s2
3s2
3s2
3p3
3p4
3p5
3p6
Figure 1
The Periodic Table
The chemical elements are gathered in a tabular way in a periodic table, which brings together the peculiar
features of each element. Several conceptually and graphically different representation of the periodic table
have been reported over the years since the first conception, by Dmitri Mendeleev. Here below a modern
representation of the periodic table showing the different "families" of the elements. All the elements of a
particular family share common features. Each line is called period and each column is called group and both
are numbered (1, 2, 3….) (Figure 2).
Figure 2
The covalent bond
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The closed shell configuration represents a particularly favorable situation and may be reached when two
neighboring atoms share electrons from their outermost shell. So, for example, when two neighboring chlorine
atoms share each one electron from their outer (valence) shell, we have a stable dichlorine molecule, featuring
a covalent bond between the two chlorine atoms.
For clarity sake, when writing a molecule structure, we normally explicit only the electrons involved in the bond
between the atoms, and, sometimes, those of the outer shell. Electrons may be symbolized either with dots or
dashes (Lewis notation). Explicit illustration of the outer shell electrons in an organic chemistry course is highly
recommended, as it allows to clearly locating the electron flow in a (ionic) chemical reaction.
1
Lower, S. http://www.chem1.com/acad/webtext/chembond/cb03.html
2
Dichlorine
electrons of the outermost sphere
shared electron pair
1 (s 2)
2 (s 2,p 6)
Cl
.. .....Cl......Cl...
Cl Cl
Cl
Cl
3 (s2 ,p 5)
Cl-Cl
simplified notations
dichlorine molecule
chlorine atom
Cl-Cl
Figure 3
The number of electrons composing the stable closed shell depends on the quantum shell considered. Period 1:
s2, period 2: 2s22p6 (octet), period 3: 3s23p6 (octet) or more 3s23p63dn (expanded octet). This (expanded octet)
situation may occur in atoms such as Si, P, S, Cl, whose d orbitals are energetically close to the most highlyoccupied s2p6 orbitals that they can become involved in electron-sharing with other atoms. As a result, they can
bond to more than four atoms, and thus need to involve more than the four pairs of electrons available in an
s2p6 octet. Some example are described here below: (Figure 4: LiH; Figure 5: PCl3), Figure 6: PCl5).
Lithium hydride
2 electrons in the outermost sphere for
each atoms 1s
Li
H
Li
H
lithium hydride
Figure 4
Phosphorous trichloride
Cl
8 electrons in the outer sphere of P in 3s and 3p orbitals (octet).
5 electrons belonging to P, of which 3 shared with 3 Cl (each Cl sharing
in turn one electron) plus 2 electrons as a non-bonding lone pair.
8 electrons in the outer sphere of each Cl in 3s and 3p orbitals (octet).
7 electrons belonging to each Cl of which 1 shared with P
P
P
Cl
Cl
Cl
Cl
Cl
phosphorous trichloride
P: 1s 2 // 2s2 2p6 // 3s 2 3p 3
Cl:1s 2 // 2s2 2p6 // 3s 2 3p 5
Figure 5
3
Phosphorous pentachloride
10 electrons in the outer sphere of P in 3s, 3p, and 3d
orbitals (expanded octet).
5 electrons belonging to P, all shared with the five Cl's.
8 electrons in the outer sphere of each Cl in 3s and 3p
orbitals (octet)
Cl
7 electrons belonging to each Cl of which 1 shared with P
Cl
Cl
Cl
Cl
P
P
Cl
Cl
Cl
Cl
phosphorous pentachloride
Cl
P: 1s 2 // 2s2 2p6 // 3s 2 3p 3
Cl:1s 2 // 2s2 2p6 // 3s 2 3p 5
Figure 6
Building a correct Lewis structure for a molecule
Knowing the above "rules" of the covalent bonding between atoms, we can now, given a raw formula of an
unknown molecule, try to play with the outer sphere electrons of the composing atoms, to locate the covalent
bonds, and therefore the connectivity ( Figure 7).
H
duet
C
N
O
F
P
S
Cl
Br
octet
octet or
expanded octet
I
Figure 7
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Accordingly, we can now build up, although in a rather raw way, the valence bonds for simple organic
molecules such as methane, ethane, ethylene, and acetylene (Figure 8).
H
1
CH4
C
4
H
H
C
H
H
H C H
H
methane
C2 H 6
2
H
H 1s 1
C 1s 2 2s 2 2p2
6
C
C 1s2 2s 2 2p 2
H
H
H 1s1
H
H
C
C
H
H
H H
H C C H
H H
H
ethane
2nd C-C bond ()
H
H
C2 H 4
4
C
2
C
H
H
H
C
C C
H
H
H
H
ethylene
H 1s 1
C 1s 2 2s2 2p2
1 st C-C bond (
2 nd C-C bond ( )
C 2H 2
2
2
C
C
H
H
C
H
H C C H
acetylene
2
2
C 1s 2s 2p
2
H 1s
1
3 rd C-C bond ( )
1 st C-C bond (
Figure 8
The shapes of the orbitals and the multiple bonds
So far we have treated atomic orbitals as undifferentiated "little balls". However, such regions of space having
the highest probability to host (up to two) electrons posses well defined geometries. More precisely, whereas s
orbitals feature a spherical geometry, p atomic orbitals are directed toward the Cartesian axes and bear a
rabbit-like shape (Figure
9).
z
z
x
x
y
s
z
z
y
px
x
y
py
x
y
pz
Figure 9
With this information we can now start building up very simple organic molecules, appreciate the geometry of
the atoms involved, and understand the notion of multiple bond.
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Let's take first a hydrogen atom and its lone electron in its 1s atomic orbital (Figure
10, top). Now, consider a
carbon atom. Its outermost 2nd quantum shell contains four electrons: two electrons in the 2s orbital and two
electrons in two singly occupied p orbitals (Figure
10,
bottom, A). Now, imagine to promote one of the
electrons of the doubly filled 2s level into the empty p orbital (Figure
10,
bottom, B). The appropriate
combination of these four singly occupied orbitals (1s and 3p) gives rise to four identical orbitals directed along
a tetrahedron: the so called sp3 hybridation (Figure 10, bottom, C).
H
1s1
C
2s1
2p2
2s2
A
2p 3
4e - sp 3
B
C
Figure 10
If we now take four of the above described hydrogen atoms and move them toward the fours lobes of the sp3
hybrid of the carbon atom we "create" fours C-H bonds. We have "created" methane!
Figure 11.
H
+
1s
1
1s1
1s
1
1s
H
H
H
1
4e- sp3
Figure 11
Let's now start again from a carbon atom, and proceed as in Figure 10 until stage B. However, now we combine
in the hybridization process just three of the fours orbitals, leaving the p z atomic orbital untouched, we create
an sp2 hybrid for carbon atom (Figure 12). Such a hybrid has a trigonal (planar) geometry.
z
z
x
C
y
y
2p2
2s 2
z
x
2s 1
A
2p3
B
4e - sp2
C
Figure 12
If we now combine two such C atoms via interpenetration of two sp2 lobes and "add" four hydrogen atoms to
the remaining four sp2 lobes we are on the edge of creating ethylene. However, for the moment we have built
only the single C-C bond. Interaction between the two unused pz orbitals of each C atom generates a second CC bond. Thus, ethylene features a double C=C bond. As we can see, the two carbon bonds are different. The
first one is the result of a combination of the sp2 hybridized lobes, whereas the second results from the
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interaction of the two parallel pz atomic orbitals. The former CC bond is stronger and is defined a - bond,
whereas the latter is weaker and is known as a -bond (Figure 13).2
-bond
H
H
C C
H
H
-bond
Figure 13
If we now repeat the process of Figure 10 until stage B combining in the hybridization process just one s and
and one p, we create an sp hybrid for carbon atom (Figure 14). Such a hybrid has a linear geometry. In this
case the remaining unused p orbitals are two; py and pz.
z
z
x
x
C
y
y
2s 1
2p2
2s 2
z
A
y
4e - sp
2p3
B
C
Figure 14
Now, combining two such C atoms via their sp lobes and adding two hydrogen atoms to the remaining two sp
lobes we can create the σ bonds of acetylene. Interaction between the two unused pz orbitals of each C atom
and, similarly, between the two unused py orbitals, generates two mutually orthogonal π-type CC bonds. Thus,
acetylene has a CC triple bond (Figure 15)!
-bond
z
z
H C C H
y
y
-bond
Figure 15
2
A -bond is a bond whose orbital does not change of sign upon rotation of 180° around the bond axis,
whereas a -bond does.
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