Electron Configuration of the Elements:
In order to understand organic chemistry we must learn the electron configuration of the first
20 elements (H to Ca) plus Br and I.
Electrons are continuously buzzing around the nucleus at mind-boggling speeds (ca. 1/10 the
speed of light). We don’t know the exact position of electrons from one moment to the next
(Heizenburg uncertainty principle) but we do know that their movement is not entirely
random.
Electrons fly within well-defined flight paths (orbitals) around the nucleus. Each orbital can
hold a maximum of 2 electrons. Think about the heavier elements on the periodic table, with
100+ electrons flying around the nucleus in 50+ different flight paths. Inevitably, some of
the orbitals overlap. Just imagine how busy their flight controllers must be while trying to
prevent all those flying electrons from colliding.
The orbitals lowest to ground zero (the nucleus) are lowest in energy and are occupied by
electrons before the outer, high-energy orbitals. The 50+ orbitals around all atoms are
grouped into 7 different ‘energy levels’ (also called layers or ‘shells’) with n = 1 being the
shell closest to the nucleus (lowest energy) and n = 7 being the farthest from the nucleus
(highest energy). Study the order of orbitals and shells in the planetary model of the atom.
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PLANETARY MODEL OF THE ATOM SHOWING ENERGY LEVELS n = 1 to 6
P shell (n = 6)
6s  1, 6p  3, 6d  5, 6f  7, 6g  9, 6h  11
O shell (n = 5)
5s  1, 5p  3, 5d  5, 5f  7, 5g  9
N shell (n = 4)
4s  1, 4p  3, 4d  5, 4f  7
M shell (n = 3)
3s  1, 3p  3, 3d  5
L shell (n = 2)
2s  1, 2p  3
K shell (n = 1)
1s  1
Nucleus
This model is not spatially correct. There is some overlap of orbitals in the 3rd shell and
higher.

Within each shell, there exist subshells or types of orbitals. The types of orbitals are
named s, p, d, f, g, h, etc.
 The 1st shell has only an s orbital, named 1s.
 The 2nd shell has both s and p-type orbitals, named 2s and 2p.
 The 3rd shell has s, p and d-type orbitals, named 3s, 3p and 3d.
 The 4th shell has s, p, d and f-type orbitals, named 4s, 4p, 4d and 4f.
 etc.
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SHAPES OF ATOMIC ORBITALS
An s orbital
(l = 0, ml = 0)
z
z
z
z
x
x
y
x
x
y
y
y
A px orbital
A py orbital
A pz orbital
(on the x axis)
(l = 1, ml = -1)
(on the y axis)
(l = 1, ml = 0)
(on the z axis)
(l = 1, ml = +1)
z
z
Three p orbitals
z
z
z
y
x
x
y
x
y
y
A dxy orbital
A dxz orbital
(between the x and y axes)
(l = 2, ml = -2)
(between the x and z axes)
(l = 2, ml = -1)
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x
A dyz orbital
(between the y and z axes)
(l = 2, ml = 0)
y
A dx2-y2 orbital
A dz2 orbital
(on the x and y axes)
(l = 2, ml = +1)
(on the z axis)
(l = 2, ml = +2)
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The number of each type of orbital, their shape and orientation are listed below.
 There is only one s-orbital in each shell and it is spherical.
 There are three p-orbitals in the 2nd and all higher shells. The three p-orbitals
are propeller shaped and are oriented along an x, y or z axis in space. They are
named px, py and pz, respectively.
 There are five d-orbitals in the 3rd and all higher shells. Four of the five
d-orbitals look like four-leaf clovers each oriented differently around the
nucleus. The fifth d-orbital looks like a propeller inside a donut. They are
named dxy, dxz, dyz, dx2-y2 and dz2.
 There are seven f-orbitals in the 4th and all higher shells. All but one have six
lobes. Each one is oriented differently around the nucleus.
In writing the electron configuration of the elements we fill lowest energy orbitals first
(Aufbau principle), with a maximum of 2 electrons per orbital –with opposite spins (Pauli
Exclusion principle). Orbitals of the same energy level (‘degenerate orbitals’) are all singly
filled (half-filled) before electrons pair up. This occurs, for example in the 2px, 2py, and 2pz
orbitals.
The filling order (increasing energy level) of the various orbitals is shown in the following
chart.
ns
(n-2)f
(n-1)d
np
H
1s
He
Li
2s
2p
Ne
Na
3s
3p
Ar
K
4s
3d
4p
Kr
Rb
5s
4d
5p
Xe
Cs
6s
4f
5d
6p
Rn
Fr
7s
5f
6d
7p
Uuo
Compare this table with a periodic
table. The filling order is the same as
the layout of the s-, p-, d- and fblocks on the periodic table, i.e., this
is the filling order. Note the patterns.
The orbital filling order is:
ns, (n-2)f, (n-1)d, np.
The p-orbitals begin filling after the
2s orbital.
The d-orbitals begin filling after the
4s orbital.
The f-orbitals begin filling after the
6s orbital.
Reading the table left to right and top to bottom, the orbital filling order is as follows:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p
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For the purpose of learning organic chemistry, we need only study the electron configuration
of the first 20 elements (H to Ca) plus Br and I, i.e., in the following orbitals:
1s 2s 2p 3s 3p 4s
Br and I have electrons in the 4p and 5p orbitals respectively, but their electron arrangement
is analogous to F and Cl.
The electron configuration of atoms is shown using a notation in which the number of
electrons in each orbital is written as a superscript. The orbital is shown as a line, _ or as a
circle, O. Each electron in the orbital is written as an arrow, . The direction of the arrow is
either up, , (indicating clockwise rotation) or down,  , (indicating counterclockwise
rotation). Complete the following table.
Full Orbital Notation
1s
2px
2py
2pz
__
1s1
He __
1s2
1
2
2s
Simplified
Notation
H
3
Li
4
Be
5
B
__
__
__
__
__
__
__
__
__
6
1s
__
__
2s
__
__
2px
__
__
2py
__
__
2pz
__
__
__
__
__
__
__
__
__
__
__
__
__
__
__
__
__
__
__
__
__
__
C
7
N
8
O
9
F
10
Ne
11
Na
1s
Write out the full electron
configuration of carbon:
6C
1s2
1s2
1s2
2s1
2s2
2s2
2p1
1s2
1s2
2s2
2s2
2p2
2p3
1s2
1s2
1s2
1s2
2s2
2s2
2s2
2s2
2p4
2p5
2p6
2p6
3s
__
2s
2px 2py 2pz
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ENERGY
The 2px, 2py, & 2pz
orbitals are equal in
energy ('degenerate').
outer, valence electrons
bonding electrons
Problem: Write out the electron configuration for Mg through Ca in both 'Full Orbital
Notation' and 'Simplified Orbital Notation'. Recall that the 4s orbital is filled before the
3d orbital.
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Full Orbital Notation
3s 3px 3py 3pz
12
M
g
13

[Ne]
3s2
Si
15
P
16
S
17
Cl
Ar
19
[
4s
Al
14
18
[Ne]
Simplified Orbital Notation
K
Ca
The outermost occupied shell is referred to as the ‘valence’ shell. Orbitals of the valence
shell are thus ‘valence orbitals’ and electrons in the valence orbitals are ‘valence electrons’.
The outer (valence) electrons are transferred or shared in chemical reactions. Chemistry is
understood in terms of valence electron arrangement.
The number of valence electrons determines the ability of an atom to combine with other atoms
 The number of covalent bonds an atom forms to become isoelectronic with its nearest
noble gas is called its ‘covalence’. (Isoelectronic means ‘having the same valence
electronic configuration’.)
 The number of valence electrons in an atom is shown with a Lewis Symbol. One dot is
drawn for each valence electron. The dots are placed into four positions (one for each of
the one s plus three p orbitals) around the symbol of the element, i.e., north, south, east or
west. Once all four positions are singly filled, electrons (dots) are paired up until a
maximum of 8 valence electrons (dots) have been drawn.
some Lewis symbols
.
.
B
.
.
.C.
.
..
. N.
.
Although Lewis symbols do not always show the lowest energy electron arrangement of an
unbonded atom, they are a good depiction of the electron arrangement just prior to bonding.
When dots (electrons) are drawn, each orbital is first half-filled before electrons are paired up
in orbitals.
The arrangement of elements in the periodic table is based on the number of valence
electrons. For example, elements in Group IVA have 4 valence electrons.
For all representative elements (A-group elements), the number of valence electrons equals
the group number.
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 Complete the following table. Note that He is an exception. Although it has 2 valence
electrons like Be and Mg, it is unreactive (like other noble gases) and therefore has a
valence of 0.
Group #
1A
2A
3A
4A
5A
6A
7A
8A
# valence
electrons
1
2
3
4
5
6
7
8
covalence
(# bonds)
Period
1
H
Period
2
Li
Be
B
C
N
O
F
Ne
Period
3
Na
Mg
Al
Si
P
S
Cl
Ar
Period
4
K
Ca
Br
Kr
electron
config.
s1
s2
show lone
pairs of
electrons
after
bonding
H
He
Be
s2 p1
B
s2 p2
C
s2 p3
N
s2 p4
O
I
Xe
s2 p5
s2 p6
F
 Note that covalence is the same as the group number for groups 1A to 4A, but covalence
is equal to [8 - (group number)] for groups 5A to 8A.
 It is very important to appreciate the relationship between the Lewis symbols and the
number of covalent bonds formed by an atom (its covalence). Usually, in most stable
organic compounds, the atoms form a covalent bond for each unpaired electron in the
Lewis symbol of the atom.
Study the bonding arrangements of the neutral atoms shown below. Note that all single
(unpaired) electrons in a Lewis structure will bond (as shown in the bonded structure). The non
bonded electron pairs (‘lone pairs’) may either remain unbonded or form two additional bonds
per electron pair. The two additional bonds may be two single bonds or one double bond
Octet Rule:
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Note however that 2nd period elements (B, C, N, O and F) will never have more than 4 bonds
(8 electrons) around themselves as they can only use four orbitals - their 2s and 2p orbitals
for bonding (‘octet rule’).
Hypervalent Atoms:
3rd period elements and higher (Si, P, S, Cl and Br) can form more than 4 bonds (more than 8
electrons) by using their d-orbitals. Examples include PCl5, SF6, ClF7, and BrF7. Such
elements that exceed the octet rule are called ‘hypervalent’.
Group #
Lewis
Symbol
3A
.
.
B
.
4A
.
.C.
.
5A
.
B
.
N
.
Al .
.
. Si .
.
..
..
: F.
..
..
..
:F
O
..
..
N
C
.
..
P
.
..
P
Bonded
Structure
..
N
.
..
. ..O .
O
..
..
C
C
Lewis
Symbol
N.
7A
..
C
Bonded
Structure
..
6A
.
..
..
..
. ..S .
: Cl .
..
: Cl
..
..
S
..
Si
: Cl
..
Al
P
S
..
Si
Cl
..
P
S
Cl
Carbon and nitrogen are the only two elements that can form a triple bond.
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