Electron Configuration of the
Elements
Hydrogen Emission Spectrum
When hydrogen gas (H2) is placed in a CRT
and a high voltage electrical current
passed through it, the tube glows a violet
colour.
Johann Balmer 
When viewed through a spectroscope
(prism), we observe four discrete lines—
and NOT a continuous spectrum:
When an electron in the ground state of H
absorbs energy, it gets promoted into a
higher energy level.
The electron is unstable in this higher
energy level.
When the electron falls back to the ground
state, energy is given off.
This explains the “bands” of light emitted
from a hydrogen discharge tube.
Here’s another way to look at it:
Hydrogen Emission Spectrum
Electrons can only exist in certain energy
levels (n)
n = 1, n = 2, n = 3, n = 4, etc
Energy levels in atom are
quantized.
This means that only certain E levels are
allowed.
• Each E level has one or more sublevels
called orbitals
• An orbital is a region of space
where there is a high probability
of finding an electron
• Each orbital can hold a maximum of two
electrons.
• Electrons in an orbital will have opposite
spin, designated ↑ (clockwise spin)
or ↓ (counterclockwise spin).
high “probability” ?
Heisenberg Uncertainty Principle
We cannot simultaneously know the position
and the momentum of an electron
Back to orbitals . . .
• For n = 1 there is only one sublevel, called
an s orbital.
• Since this orbital is in the first energy level,
it is called a 1s orbital.
• s orbitals are spherical.
•
•
•
•
•
For n = 2 there are two sublevels:
2s orbital (one of these)
2p orbital (three of these)
a p orbital looks like this
A set of three p orbitals looks like this
• We refer to the individual p orbitals as
px, py, pz.
Let’s put these orbitals together . . .
• For n = 3 (the third energy level) there are
three sublevels:
3s orbital (one of these)
3p orbital (three of these)
3d orbital (five of these—see next slide)
NB. Each orbital holds a maximum of 2
electrons
The d-orbitals
A funky look at d-orbitals
Your “bottom line” with d-orbitals:
• There are five of them in each set.
• eg. there are five 3d orbitals; five 4d
orbitals, etc
• 2 electrons in each, for a maximum of
10 electrons
• How many columns are in the Transition
Metal block (d-block) in the periodic table?
10 columns in the transition metals (5x2).
For n = 4 (the fourth energy level) there
are four sublevels:
4s orbital (one of these)
4p orbital (three of these)
4d orbital (five of these)
4f orbital
(seven of these—see next slide)
f-orbitals
How do electrons
fill orbitals?
Aufbau Principle
• aka “Building-up” Principle
• Electrons occupy orbitals beginning from the
lowest energy orbital
(i.e. the orbital closest to the nucleus)
Start by filling 1s orbital
• How many electrons per orbital?
Each orbital can hold a maximum of two
electrons—of opposite spin, don’t forget
Here is the order in which orbitals
are filled . . .
Note the peculiarity . . .
3s is followed by
3p, which is followed by
4s, which is followed by
3d.
There are others . . .
(help is on the way)
How do the electrons of 7N fill the orbitals?
1s2
2s2
2p3
Overall for 7N: 1s2 2s2 2p3
Hund’s Rule
More stable than . . .
Hund’s Rule
• When filling p, d, f orbitals, pair electrons
only when necessary
Aufbau Principle Mnemonic Device
Let’s write some electron
configurations . . .
1H
2He
3Li
4Be
5B
6C
↓
↓
1s1
1s2
1s2 2s1
1s2 2s2
1s2 2s2 2p1
1s2 2s2 2p2
2 2s2 2p6
Ne
1s
10
11Na
2 2s2 2p6 3s1
Na
1s
11
2 2s2 2p6 3s2
Mg
1s
12
2 2s2 2p6 3s2 3p1
Al
1s
13
↓
↓
2 2s2 2p6 3s2 3p6
Ar
1s
18
2 2s2 2p6 3s2 3p64s1
K
1s
19
2 2s2 2p6 3s2 3p64s2
Ca
1s
20
2 2s2 2p6 3s2 3p64s2 3d1
Sc
1s
21
In the Periodic Table, what is the connection
between the outermost electron
configuration and family (column)?
Alkali metals
end in s1
Alkali earth metals
end in s2
Halogens
end in p5
Noble Gases
end in p6
Groups (families) in PT
Putting it all together . . .
• To write the electron configuration of any
element, use the periodic table (play
“Battleship”) and the Aufbau Principle
mnemonic device.
Write the complete electron configuration for
1s2 2s2 2p6 3s2 3p64s2 3d10 4p6 5s2 4d10 5p1
____________________________________
[Kr]
preceding noble gas only
2 4d10 5p1
In
[Kr]
5s
49
Practice
Write the electron configuration for each of
the following
15P
2 2s2 2p6 3s2 3p3
P
1s
15
33As [use noble gas core abbreviated form]
2 3d10 4p3
As
[Ar]
4s
33
more practice . . .
13Al
2 2s2 2p6 3s2 3p1
Al
1s
13
26Fe
2 2s2 2p6 3s2 3p6 4s2 3d6
Fe
1s
26
44Ru [ ]
2 4d6
Ru
[Kr]
5s
44
52Te [ ]
2 4d10 5p4
Te
[Kr]
5s
52
Exceptional Electron
Configurations
1. Write the expected electron configuration
of 24Cr
[Ar]4s2 3d4
Actual electron configuration is
[Ar]4s1 3d5
Special stability associated with
half-filled p, d, f orbitals
Now write the electron configuration for 42Mo
14d5
Mo
[Kr]
5s
42
Notice any similarity with Cr?
2. Write the expected electron configuration
of
29
[Ar]4s2 3d9
Cu:
Actual electron configuration is
[Ar]4s1 3d10
In this way Cu has completely filled 3rd
energy level
(Copper is a very stable metal)
Now write the electron configuration for
silver (47Ag) and gold (79Au). Use the
noble gas core abbreviated forms.
1 4d10
Ag
[Kr]
5s
47
1 4f14 5d10
Au
[Xe]
6s
79