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Chapter 4
Electron Configurations and
Quantum Chemistry
Electron configurations determine
how an atom behaves in bonding
with other atoms!
Topics rearranged from your text, pages 90-116.
Atomic Emissions/Abortions removed
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Anyone who says that they can contemplate quantum
mechanics without becoming dizzy has not understood
the concept in the least.
-Niels Bohr Slide 1
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The Bohr Model
• Niels Bohr
– rebuilt the model of the atom placing the electrons
in energy levels.
• Quantum chemistry
– a discipline that states that energy can be given off
in small packets or quanta of specific size.
• What would happen to an electron if the right
sized quanta of energy was added to it?
EXCITED STATE
• What would happen when the electron came
back down to its ground state? Ground state
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Slide 2
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Electron Configurations - overview
• Bohr model
– electrons exist in specific energy levels.
• Electron orbitals (shapes)
– Within each energy level, the orbits the electrons
can occupy.
• Within each orbital
– electrons can be set “spin up” or “spin down”
• Electron configuration
– The configuration of electrons in their levels,
orbitals, and spins.
• Modern Quantum Model
– Electron exists in electron configurations
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Slide 9
Energy Levels (n)
• The electrons exist in energy
levels or shells.
• The first energy shell can
hold only 2 electrons.
Old School: Back
“KLM notation”
2
– Hydrogen and Helium in their
ground state have electrons
that occupy this shell.
8 18
32
• The second shell can hold 8 Shells
electrons.
All shells after
• The third can hold 18
three can hold
32 electrons.
electrons.
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Slide 10
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Orbitals (Shapes)
• Orbitals
– electrons travel in set paths.
– These paths form shapes, called
orbitals.
• Each “shape” can hold 2
electrons
• The smallest orbital is the “s”
orbital. The “s” orbital:
– Has only 1 shape (holds 2 e-)
– Is spherical in shape
– Is the lowest energy orbital
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s-2
Slide 11
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p-Orbitals
• The 2nd orbital shape is the “p” orbital shape.
• There are 3 “p” shapes, each holding 2
electrons, for a total of 6 electrons in the “p”
orbitals.
• The “p” orbitals are:
– Dumbbell-shaped
s-2
– Higher in energy than the “s”
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p-6
Slide 12
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d-Orbitals
• The 3rd orbital shape is the “d” orbital shape.
• There are 5 “d” orbital shapes, for a total of 10
electrons in the “d” orbitals.
• “d” orbitals are higher in energy than “p”
orbitals.
s-2
p-6
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d-10
Slide 13
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f-Orbitals
• The last orbital
shape is the “ f ”
orbital shape.
Electrons in f orbitals
are very high in energy
– “ f ” orbitals have
irregular shapes
due to quantum
tunneling.
– There are 7 “ f ”
shapes, for a total
of 14 electrons.
s-2
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p-6
d-10
f-14
Slide 14
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“Blocks” of the periodic table…
• The periodic table tells us in which orbital the
last electron should be found.
– The last electron in an atom is found in the…
p orbitals
s orbitals
d orbitals
f orbitals
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Slide 15
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Electron “Spin”
• Electrons can be “spin up” or “spin down.”
– (by convention, an electron that is alone is “spin up”)
• Hund’s Rule
– As electrons fill orbitals, they first fill each shape
available with one electron before spin pairing.
• Pauli’s Exclusion Principle
– If two electrons share a shape, they must be spinpaired (one up and one down).
• For instance: take a “p” orbital…it has three
orbital-shapes that can hold 2 e- each.
• It would fill like this:
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Slide 16
Electron Configurations.mov
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Writing Electron Configurations
• The Aufbau principle
– electron will fill lower energy
orbitals first.
• Energy of electrons:
s Æ low energy
d Æ high energy
close Æ low energy
far Æ high energy
– low energy s < p < d < f high energy
– low energy nearer < farther high energy
– low energy level 1 < level 7 high energy
Total energy
• Total energy of an electron:
– Product of energy of its shell and the
energy of its orbital.
– Guess: Which is lower in energy, an
electron found in 3d or one found in 4s?
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The 4s electrons are lower in energy!
=
Shell
x
orbital shape
Slide 17
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Writing Electron Configurations
• Orbital filling diagram
– Shows how electrons fill into
levels and orbitals
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
Electron
Configurations
3d
4d 4f
5d 5f
6d 6f
7d 7f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10
6
2
14
©Bires,
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6s
5d10 6p6 7s2 5f14 6d10 7p6 6f14
5p
Å Don’t Copy this
Slide 18
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Building the Orbital Filling Diagram
• Begin by listing the shells 1, 2, 3,
1s
4, 5, 6, 7 vertically.
• These are your “s” orbitals.
2s 2p
• Next, add another column of
3s 3p 3d
number, beginning with 2.
4 s 4 p 4 d 4f
• These are your “p” orbitals.
• Do the same for “d” and “f”
5 s 5 p 5 d 5f
orbitals, beginning with “3” for
6
6
s
6
6
p
d
f
the “d” orbitals and “4” for the “f”
orbitals.
7 s 7 p 7 d 7f
• Next, add your orbital letters.
s
p
d
f
• Finally, draw diagonal lines as
shown. 1s 2 2s 2 2 p 6 3s 2 3 p 6 4s 2 3d 10 4 p 6 5s 2 4d 10 5 p 6 6s 2 4 f 14 5d 10 6 p 6 7 s 2 ...
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Slide 19
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Electron Configurations of Some Atoms
• Consider Fluorine, with 9 electrons
Notice the
position of the
last electron…
F = 1s 2 s 2 p
2
2
5
• What about Copper, with 29 electrons?
Cu = 1s 2 2 s 2 2 p 6 3s 2 3 p 6 4 s 2 3d 9
Both used
Cu = 1s 2 s 2 p 3s 3 p 3d 4 s
2
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2
6
2
6
9
2
Slide 20
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Noble Gas Shorthand
• Notice the configurations of the noble gases:
He = 1s
Ne = 1s 2 2 s 2 2 p 6 Ar = 1s 2 2 s 2 2 p 6 3s 2 3 p 6
2
• We can shorten the electron configuration of larger
elements with NGS.
2
2
6
2
Mg
=
1
s
2
s
2
p
3
s
• Consider Mg:
• We can substitute Neon’s e- config, and write Mg:
• Similarly, Titanium’s (Ti) e- config: Mg = [ Ne] 3s 2
Ti = 1s 2 s 2 p 3s 3 p 4 s 3d
2
2
6
2
• Can be shortened to:
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6
2
2
Ti = [Ar] 4s 3d
2
2
Slide 21
Ion
e
Back
configurations
• Ions (elements with more/less electrons) also have
electron configurations.
2
4
• Consider Sulfur (S):
S = [ Ne] 3s 3 p
• What if sulfur gained two electrons?
22
6
S = [ Ne] 3s 3 p
• Consider Calcium (Ca):
Ca = 1s 2s 2 p 3s 3 p 4 s
2
2
6
2
6
2
• What if calcium lost two electrons?
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Ca
2+
= 1s 2 s 2 p 3s 3 p
2
2
6
2
6
Slide 22
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Octets!
2
6
...s p ...
• Octets:
– Atoms with filled s and p orbitals in the same,
highest level.
– Have noble gas-like configurations
– Have special stability
• Both atoms and ions can have complete octets.
2-
Ne = 1s 2 2 s 2 2 p 6
S = [ Ne] 3s 3 p
Ar = 1s 2s 2 p 3s 3 p
2
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2
6
2
6
2+
2
6
Ca = 1s 2s 2 p 3s 3 p
2
2
6
2
6
Slide 23
EndBack
of
chapter 4
• Question:
– Why do the atomic radii (size) of atoms decrease as
electrons and protons are added to the atom, as
you move from left to right across a period?
• electrostatic attraction
– attraction between the electrons (-) in the shells and
the protons(+) in the nucleus – pulls the electrons in
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2002
This is what we call a periodic trend
Slide 24