South Pasadena • Honors Chemistry
Name
9 • Atomic Structure
Period
9.2
NOTES
–
Date
ELECTRON CONFIGURATION
Electron Configuration – a representation of how electrons fill atomic orbitals in an atom.
(1) Aufbau Principle – (from German aufbau meaning “building up” or “construction”) ‒ Electrons occupy the
available atomic orbitals from the lowest energy. (low to high)
(2) Pauli Exclusion Principle – (named after Austrian physicist Wolfgang Pauli) ‒ A maximum of two electrons
can occupy any atomic orbital.
(3) Hund’s Rule – (named after German physicist Friedrich Hund) ‒ If there are more than one orbital with the
same energy, occupy an empty orbital before pairing up. (if possible, don’t share)
We can use the orbital energy diagram to help us know how these orbitals are filled. A symbol “3p4” means there
are four electrons in the 3p subshell. A neutral atom of the element N has 7 electrons. Its electron configuration,
based on the rules, is: 1s2 2s2 2p3
Because the electron configuration of an element can be very cumbersome to write out, we can write it in a short
form or core notation by abbreviating the electrons that are a part of the closest previous noble gas of an atom.
For example, the short form electron configuration for a neutral atom of N is: [He] 2s2 2p3.
Example 1: Write the long- and short-form ground state electron configurations for the following species using
the orbital energy diagram:
Oxygen
(8 electrons)
Argon
(18 electrons)
3s
3p
2p
2s
Long: 1s2 2s2 2p4
Short: [He] 2s2 2p4
6 valence electrons
1s
3s
3p
Long: 1s2 2s2 2p6 3s2 3p6
Short: [Ne] 3s2 3p6
8 valence electrons
2p
2s
1s
Arsenic
4s
3s
4p
(33 electrons)
3d
Calcium ion
4s
3p
3s
2p
2s
Long: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3
Short: [Ar] 4s2 3d10 4p3
1s
5 valence electrons
(Note that 4s < 3d < 4p)
4p
(20 – 2 = 18 electrons)
3d
3p
2p
2s
Long: 1s2 2s2 2p6 3s2 3p6
Short: [Ar]
1s
0 valence electrons
(Note it’s isoelectronic with Ar)
Core Electrons – inner electrons, abbreviated by the noble gas.
Valence Electrons – electrons in the highest energy electron shell. Involved with bonding and determine chemical
properties.
When atoms lose electrons to form cations, they lose the electrons found in the highest shell. When atoms gain
electrons to form anions, they add electrons according to Aufbau Principle to the available orbital with the lowest
energy.
The Periodic Table – The order in which electrons fill is somewhat confusing. The Periodic Table can be helpful
for us to predict the electron configuration of an element because it is arranged in the way electrons are filled.
s block
1s
p block
2s
1s
2p
3s
d block
3p
4s
3d
4p
5s
4d
5p
6s
5d
6p
7s
6d
7p
f block
4f
5f
Example 2: Use the species in Example 1.
 Write the long form ground state electron configuration for the species in Example 1 using the Periodic Table
instead of the orbital energy diagram.
 In the long form electron configuration, box the electrons that are core electrons, and circle the electrons that
are valence electrons.
 Write the short form electron configuration of the species in Example 1 using the Periodic Table.
Patterns in the Periodic Table – The valence electrons of an element determine its properties. Elements in each
family across the Periodic Table have the same number of valence electrons. This explains why elements in
the same family have similar chemical properties.
Fill in the patterns observed for number of valence electrons and the expected ion formed for each family.
H
He
1s1
1s2
Li
Be
B
C
N
O
F
Ne
1
2
2
1
2
2
2
3
2
4
2
5
2s
2s
2s 2p
2s 2p
2s 2p
2s 2p
2s 2p
2p2 2p6
Na
Mg
Al
Si
P
S
Cl
Ar
1
2
2
1
2
2
2
3
2
4
2
5
3s
3s
3s 3p
3s 3p
3s 3p
3s 3p
3s 3p
3s2 3p6
General
ns1
ns2
ns2 np1
ns2 np2
ns2 np3
ns2 np4
ns2 np5
ns2 np6
Form:
# of
Valence
Electrons:
He:
1
2
3
4
5
6
7
2
Others:
8
Ion
+1
+2
+3
–3
–2
–1
0
Formed
Example 3: For each of the following elements, determine the number of valence electrons and predict the likely
charge of a typical ion that might form.
 Titanium
[Ar] 4s2 3d2
 2 valence electrons
Forms Ti2+
 Selenium
[Ar] 4s2 3d10 4p4  4 valence electrons
Forms Se2–
Notice that the elements tend to gain or lose electrons to form ions that have the same electron configuration (or
are isoelectronic) as a noble gas, so we say that they have a pseudo-noble gas configuration. Elements in the
same family tend to form ions that have the same charge because they have similar valence electron
configurations.