2. The Pauli Exclusion Principle and Electron
Configurations
The spin quantum number, ms, determines
the number of electrons that can occupy an
orbital.
ms = ±1/2
Electrons described as “spin up” or “spin down”.
An electron is specified by a set of four quantum
numbers.
3. The Pauli Exclusion Principle and Electron
Configurations
Pauli Exclusion Principle - no two electrons in
an atom may have the same set of four
quantum numbers.
Two electrons can have the same values of n, l, and
ml, but different values of ms.
Two electrons maximum per orbital.
Two electrons occupying the same orbital are spin
paired.
4. Orbital Energies and Electron Configurations
Electrons in smaller orbitals are held more
tightly and have lower energies.
Orbital size increases as the value of n increases.
True for hydrogen atoms, but not entirely true for
multielectron atoms.
As nuclear charge increases, orbital size decreases.
Electrons interact with other electrons as well as the
positively charged nucleus.
5. Orbital Energies and Electron Configurations
For electrons in larger orbitals, the charge
“felt” is a combination of the actual nuclear
charge and the offsetting charge of electrons
in lower orbitals.
The masking of the nuclear charge is called
shielding.
Shielding results in a reduced, effective nuclear
charge.
6. Orbital Energies and Electron Configurations
Effective nuclear charge allows
for understanding of the energy
differences between orbitals.
2s orbital: the small local
maximum close to the
nucleus results in an electron
with a higher effective
nuclear charge.
2p orbital: lacks the local
minimum close to the
nucleus of the 2s orbital.
Lower effective nuclear
charge for 2p electrons.
7. Orbital Energies and Electron Configurations
The energy ordering for atomic orbitals is
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f,
5d, 6p, 7s, 5f, 6d, and 7p.
An orbital’s size and penetration when treated
quantitatively produces the order of filling
represented.
Electronic configurations are written in
order of energy for atomic orbitals.
8. Hund’s Rule and the Aufbau Principle
Aufbau principle - when filling orbitals, start with the
lowest energy and proceed to the next highest energy level.
Hund’s rule - within a subshell, electrons occupy the
maximum number of orbitals possible.
Electron configurations are sometimes depicted using
boxes to represent orbitals. This depiction shows paired
and unpaired electrons explicitly.
9. 9
The Periodic Table and
Electron Configurations
The Aufbau Principle describes the electron filling
order in atoms.
10. 10
The Periodic Table and
Electron Configurations
There are two ways to remember the correct filling
order for electrons in atoms.
1. You can use this mnemonic.
11. 11
The Periodic Table and
Electron Configurations
2. Or you can use the periodic chart .
12. Example Problem 6.6
What is the electron configuration for the
sulfur atom?
1s2 2s2 2p6 3s2 3p4
13. Hund’s Rule and the Aufbau Principle
A simplified depiction uses superscripts to indicate
the number of electrons in an orbital set.
1s2 2s2 2p2 is the electronic configuration for carbon.
Noble gas electronic configurations are used as a
shorthand for writing electronic configurations.
Relates electronic structure to chemical bonding.
Electrons in outermost occupied orbitals give rise to
chemical reactivity of an element.
[He] 2s2 2p2 is the shorthand for carbon
14. Hund’s Rule and the Aufbau Principle
The inner electrons, which lie closer to the nucleus,
are referred to as core electrons.
Core electrons can be represented by the noble gas with
the same electronic configuration.
The outer electrons are usually referred to as
valence electrons.
Valence electrons are shown explicitly when a noble gas
shorthand is used to write electronic configurations.
Valence electrons determine reactivity.
15. Example Problem 6.7
Rewrite the electron configuration for sulfur
using the shorthand notation.
[Ne] 1s2 2s2 2p6
16. The Periodic Table and Electron Configurations
The periodic table and the electronic
configurations predicted by quantum mechanics
are related.
The periodic table is broken into s, p, d, and f blocks.
Elements in each block have the same subshell for the
highest electron.
Structure of periodic table can be used to predict
electronic configurations.
17. The Periodic Table and Electron Configurations
The shape of the periodic table can be broken down into blocks
according to the type of orbital occupied by the highest energy
electron in the ground state.
We find the element of interest in the periodic table and write its
core electrons using the shorthand notation with the previous
rare gas element. Then we determine the valence electrons by
noting where the element sits within its own period in the table.
18. Example Problem 6.8
Use the periodic table to determine the
electron configuration of tungsten (W),
which is used in the filaments of most
incandescent lights.
[Xe] 6s2 4f14 5d4
19. 19
The Periodic Table and
Electron Configurations
Now we will use the Aufbau Principle to
determine the electronic configurations of the
elements on the periodic chart.
1st row elements.
2
2
1
1
1sHe
1sH
ionConfigurat1s
1
1 1sH
ionConfigurat1s
20. 20
The Periodic Table and
Electron Configurations
2nd row elements.
•Hund’s rule tells us that the electrons will fill the
p orbitals by placing electrons in each orbital
singly and with same spin until half-filled. Then
the electrons will pair to finish the p orbitals.
27. 27
The Periodic Table and
Electron Configurations
1102
31 4p3d4sArArGa
ionConfigurat4p4s3d
28. 28
The Periodic Table and
Electron Configurations
Now we can write a complete set of quantum
numbers for all of the electrons in sodium (for
example).
Na
11Na.
When completed there must be one set of 4 quantum
numbers for each of the 11 electrons in
(remember Ne has 10 electrons)
1
11 s3NeNeNa
ionConfigurat3p3s