• Erwin Schrodinger proved
the idea of the “Planetary
Atom” wrong
• He said the e- does not move
around in fixed orbits
• He proved that the e- moves
in 3D areas around the
nucleus
• He made a complicated
math equation describing
the area of an e-
• BUT…its just a probability of
where an e- may be in the
atom
• We do not know exactly
where an e- is in an atom
• e- have different energy
levels called principal
quantum numbers (n)
• Within the energy levels
are sublevels that have their
own significant shape
• 3D non-circular areas where
e- can be
• Inside sublevels there are
orbitals, where e- are
arranged
• s is the first and
simplest sublevel
• Has only 1 orbital that
holds 2 electrons
• p is the second sublevel
• 3 orbitals
each holds 2 electrons
• A total of 6 electrons
• d is the third sublevel
• 5 orbitals
each holds 2 electrons
• A total of 10 electrons
• f is the last sublevel
• 7 orbitals
each holds 2 electrons
• A total of 14 electrons
Sublevel Orbitals Electrons
s 1 2
p 3 6
d 5 10
f 7 14
All electrons follow three
rules when filling energy
levels and sublevels. They
are:
• Aufbau Principle
• Pauli Exclusion Principle
• Hund’s Rule
• Each electron occupies
the lowest energy
orbital available
Here’s how to remember…
• Principal quantum numbers are like floors of a hotel
• e- enters the lowest level first
The Hotel Californium
1st floor
2nd floor
5th floor
n = 1
n =2
n = 5
= electron
• A maximum of two
electrons may occupy a
single orbital, but only
if the electrons have
opposite spins
Here’s how to remember…
Spin one index finger away
from your
Spin the other index finger
toward you
• Single electrons with
the same spin must
occupy each equal-
energy orbital before
additional electrons
with opposite spins can
occupy the same
orbital.
Here’s how to remember…
“Strangers on a bus”
• Do you sit next to a stranger or pick an empty seat?
• Pick an empty seat…unless there isn’t one, then you sit with the stranger
• Includes a box for each
of the atom’s orbitals
• Boxes are filled with
arrows representing
electrons
Up and down facing
arrows to show opposite
spins
Orbital Diagrams
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f • Start from bottom of
Aufbau diagram.
• Remember orbital table
for the sublevels
1s2s2p3s3p4s3d
Incr
easi
ng e
ner
gy
• Remember the rules
electrons follow when
filling energy levels,
sublevels and orbitals
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule
• Let’s determine the
orbital diagram for
Phosphorus
15 electrons
Filling in Orbital Diagrams:
• The first 2 e- go into 1s
orbital (line )
• only 13 more to go...
• Next e- go into 2s orbital
• only 11 more…
• Next e- go into 2p orbital
• only 5 more…
• Next e- go into 3s orbital
• only 3 more…
• Last e- go into 3p orbital
• e- by themselves before
being paired
• 3 unpaired e-
1s2s2p3s3p4s3d
Notice the opposite
spins (one ↑, one ↓)
• Determine the orbital
diagram for the
following:
Cobalt (Co)
Germanium (Ge)
1s2s2p3s3p4s3d
Cobalt – 27 e- Germanium – 32 e-
1s2s2p3s3p4s3d
4p
Electron
Configurations• 1s2 2e-
• 1s22s2 4e-
• 1s22s22p63s2 12e-
• 1s22s22p63s23p64s2 20e-
• 1s22s22p63s23p64s23d10
4p65s2 38e-
• 1s22s22p63s23p64s23d10
4p65s24d105p66s2 56e-1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
Example: Oxygen
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f• 1s2
• 2 electrons
• 4 electrons
• Total 8 electrons
• Stop once the # e- = the
atomic number
• Last sublevel doesn’t have
to be completely filled
8
16O
8 electrons2s2 2p4
• Use your periodic
tables to find the
atomic numbers and
the electron
configurations of the
following elements:
Silicon
Uranium
• Silicon
1s22s22p63s23p2
• Uranium
1s22s22p63s23p64s23d10
4p65s24d105p66s24f14
5d106p67s25f4
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