This document discusses electron configuration and how it is described using quantum numbers. It explains the main concepts like energy shells (n), subshells (s, p, d, f), orbitals and how electrons fill these according to the Aufbau principle and Hund's rule. Examples are provided to show the electron configurations of elements like hydrogen, helium, lithium and how the configurations change as the atomic number increases. Practice problems are included at the end to determine configurations of additional elements.

1. Electron
Configuration
IB Chemistry Power Points
Topic 2
Atomic Structure
www.pedagogics.ca

2. Going beyond 2,8,8,2
Electron configuration in atoms can be described by
terms called quantum numbers – no two electrons can
have the same quantum number!
st
1
Term: shell (n)
- main energy level
n = 2
n = 1
n = 3
lone electron
of hydrogen
energy
Issue with this graphic?

3. 2
nd
Term: subshell
- designated by s, p, d, f
- designates the sub-energy level
within the shell.
- refers to the shape(s) of the
volume of space where
n = 3
n = 2
1s electrons are be located.
The first shell (1) has one subshell (s).
The s subshell has 1 spherical shaped orbital
orbitals are volumes of space where the
probability of finding an electron is high
energy

4. The Electronic Configuration of Hydrogen
1s
Hydrogen has one electron located in the first shell.
(Aufbau principle – electrons will occupy the lowest
energy orbitals first)
The first shell has only one subshell (s).
The s subshell holds a single s orbital.
Electronic configuration
1
1s
shell
subshell
# of electrons present
energy
1s 
Orbital Energy Level Diagram

5. The Electronic Configuration of Helium
He: Atomic # of 2, 2 electrons in a neutral He atom
1
H 1s
2
He 1s
1s He 1s 
energy
the maximum number of electrons in an orbital is TWO
if there are 2 electrons in the same orbital they must have an
opposite spin.
This is called Pauli’s Exclusion Principle

6. Lithium (Li)
Li: Z=3 Li has 3 electrons.
1s
2
nd
shell
1s
The 2nd shell (n= 2) has 2
subshells which are s and p.
The s subshell fills first!
(Aufbau Principle)
2s
2p
2s 
Li 1s 
Orbital Energy Level Diagram
2
2s
Li 1s
1
Electronic configuration
energy

7. Subshells so far
- designated by s, and p
- refers to the shape(s) of
the volume in which the electron
can be located.
- also designates an energy level
within the shell.
- relative energy: s < p
s subshell: spherical
1 orbital
x
y
z
x y z
p subshell: pair of lobes, 3 orbitals, each holds 2 electrons

8. Berylium (Be)
Be: Z=4 Be has 4 electrons.
Be 1s
2
2s
2 2s 

Be 1s 
Electronic configuration
Orbital Energy Level Diagram
Boron (B) has 5 electrons, the s subshell is full so the 5
1s
2
nd
shell
2s
2p
B 1s
2
2s
2
2p
1
2p 
2s 
B 1s 
th
electron occupies the first orbital in the p subshell
energy

9. Carbon (C)
C: Z=6 C has 6 electrons.
1s
2
nd
shell
2s
2p C 1s
2
2s
2
2px
1
py
1
2p  
2s 
C 1s 
C 1s
2
2s
2
2p
2
The 6
th
electron occupies an
empty p orbital. This illustrates
“Hund’s Rule” – electrons do not
pair in orbitals until each orbital
is occupied with a single electron.
The electron configuration is
But always written as

10. 2p   
2s 
N 1s  1s
2
2s
2
2p
3
2p   
2s 
O 1s  1s
2
2s
2
2p
4
2p   
2s 
Ne 1s  1s
2
2s
2
2p
6

11. Practice
Use the sheets provided to fill out orbital diagrams and
determine the electron configuration for the following
elements
1. Fluorine
2. 56Fe
3. Magnesium - 22
4. 131I
5. Potassium – 42
6. 75Ge
7. Zirconium – 90
8. 41Ca2+

12. Practice
Use the sheets provided to fill out orbital diagrams and
determine the electron configuration for the following
elements
1. Fluorine 1s
2
2s
2
p
5
2.
56
Fe 1s
2
2s
2
p
6
3s
2
3p
6
4s
2
3d
6
3. Magnesium – 22 1s
2
2s
2
p
6
3s
2
4.
131
I 1s
2
2s
2
p
6
3s
2
3p
6
3d
10
4s
2
4p
6
4d
10
5s
2
5p
5
5. Potassium – 42 1s
2
2s
2
p
6
3s
2
3p
6
4s
1
6.
75
Ge 1s
2
2s
2
p
6
3s
2
3p
6
4s
2
3d
10
4p
2
7. Zirconium – 90 1s
2
2s
2
p
6
3s
2
3p
6
4s
2
3d
10
4p
6
5s
2
4d
2
8.
41
Ca
2+
1s
2
2s
2
p
6
3s
2
3p
6

13. Electron Configurations and the Periodic Table
So far, we have seen how the subshell model provides
and explanation for the patterns in ionization energy
we see in the periodic table.
You have also seen how to write electron configurations
Example CALCIUM  1s
2
2s
2
2p
6
3s
2
3p
6
4s
2
-
Principle energy level subshell # of e
Calcium can also be written shorthand as:
2
[Ar]4s

14. The organization of the Periodic table correlates directly to
electron structure

15. Condensed electron configurations – for example the electron
2
10
5
configuration of bromine can be written [Ar] 4s
3d
4p
Read questions carefully – many IB questions require you
to write the FULL electron configuration

16. You are responsible for configurations up to Z = 36 (Kr). The table
works well for this with the exception of Cr and Cu

17. Chromium’s configuration is:
1
3d
[Ar]4s
5
Copper’s configuration is:
[Ar]4s
1
3d
10
These configurations are energetically more stable
than the expected arrangements. KNOW THEM!

18. Electron configuration of ions:
In general, electrons will be removed from orbitals (ionization) in the
reverse order that the orbitals were filled. In other words, electrons
vacate higher energy orbitals first.
The exception: TRANSITION METAL IONS
When these ions form, electrons are removed from the valence shell s
orbitals before they are removed from valence d orbitals when transition
metals are ionized.
For example: Cobalt has the configuration [Ar] 4s
2
3d
7
The Co
2+
3+
and Co
ions have the following electron configurations.
Co
2+
: [Ar] 3d
7
Co
3+
: [Ar] 3d
6

19. Condensed electron configurations – for example the electron
2
10
5
configuration of bromine can be written [Ar] 4s
3d
4p
1. Si ___________________________
2. S2- ___________________________
3. Rb+ ___________________________
4. Se ___________________________
5. Ar ___________________________
6. Nb ___________________________
7. Zn2+ ___________________________
8. Cd ___________________________
9. Sb ___________________________

20. Review: the principles involved
Aufbau Principle: electrons will fill the lowest energy orbitals
first
Hund’s Rule: the most stable arrangement of electrons in
orbitals of equal energy is where there is the maximum
number of unpaired electrons all with the same spin.
Pauli’s Exclusion Principle: A maximum of two electrons can
occupy a single orbital. These electrons will have opposite
spins.