How do the Energy Levels Fill Up with Electrons?


Recall:
◦ Werner Heisenberg formulated the Uncertanity Principle
that states it is impossible for us to know an electron’s
exact position (where it is) and momentum (where it is
going)
◦ As a result, we cannot identify specific orbits that
electrons travel in
◦ We can only identify regions of space within an atom
where an electron is most likely to be found
 ORBITALS!
◦ Schrodinger’s complex math equation allows us to:
 Calculate the shape of the electron cloud
 Probability of finding the electron at distinct locations within
those clouds
An Introduction to Electron
Configurations

Complete the activity “Welcome to Atomos
Apartments!” on page 208

We use electron configurations
◦ The way electrons are arranged in atoms

There are rules to follow!
◦ Aufbau principle
 Electrons are added one at a time to the lowest
energy orbitals available until all the electrons
of the atom have been accounted for
 “aufbau”
 German for ‘build up or construct’
1s
2s 2p
3s 3p 3d
4s 4p 4d
4f
5s 5p 5d
5f

Pauli’s Exclusion Principle
◦ An orbital can hold only two electrons

Hund’s Rule
◦ “Electrons must fill a sub-level such that each
orbital has a spin up electron before they are paired
with spin down electrons”

A bus analogy:
◦ If you enter a bus and don’t know anyone on it, you will
pick a seat that is completely empty rather than one that
already has a person in it




Electrons fill in order from lowest to highest energy
The Pauli exclusion principle holds. An orbital can hold
only two electrons
Two electrons in the same orbital must have opposite
signs (spins)
You must know how many electrons can be held by
each orbital
◦
◦
◦
◦

2 for s
6 for p
10 for d
14 for f
Hund’s rule applies. The lowest energy configuration
for an atom is the one having the maximum number of
unpaired electrons for a set of orbitals
◦ By convention, all unpaired electrons are represented as
having parallel spins with the spin “up”

Just a thought…
◦ How do you determine the number of electrons in
an element?

Examples:
◦
◦
◦
◦
Oxygen
Magnesium
Argon
Scandium

Use the Noble Gas symbol to abbreviate or
shorten the electron configuration
◦ Krypton
◦ Rubidium
◦ Zirconium
Use Quantum Numbers!
Each electron has a specific ‘address’ in the
space around a nucleus

An electrons ‘address’ is given as a set of
four quantum numbers

Each quantum number provides specific
information on the electrons location

state
town
house number
street

state (energy level) - quantum number n

town (sub-level) - quantum number l

street (orbital) - quantum number ml

house number (electron spin) - quantum
number ms



Same as Bohr’s n
Integral values: 1, 2, 3, ….
Indicates probable distance from the nucleus
◦ Higher numbers = greater distance from nucleus
◦ Greater distance = less tightly bound = higher
energy


Integral values from 0 to n - 1 for each
principal quantum number n
Indicates the shape of the atomic
orbitals
Table 7.1 Angular momentum quantum
numbers and corresponding atomic
orbital numbers
Value of l
0
1
2
3
4
Letter
used
s
p
d
f
g


Integral values from l to -l, including zero
Relates to the orientation of the orbital in
space relative to the other orbitals
◦ 3-D orientation of each orbital

An orbital can hold only two electrons, and
they must have opposite spins
◦ Spin can have two values, +1/2 and -1/2

Pauli Exclusion Principle (Wolfgang Pauli)
◦ "In a given atom no two electrons can have the
same set of four quantum numbers"

Complete the Closer on Page 206

Begin homework on page 209 – FRONT AND
BACK!