The Periodic Table!
• 8.3-5: Electron Configurations
• 8.6-8: Periodic Trends [also 9.6]
• Ion configurations and trends are
threaded throughout…
The Periodic Table holds a wealth of information
about the organization and and development of
properties, now we unlock them!
CHEM& 141 F08
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Electron Configurations
n=
1
l=
0
0
1
0
ml = 0
0
-1 0 +1
0
2
3
1
2
-1 0 +1 -2 -1 0 +1 +2
Pauli Exclusion Principle: Each individual orbital takes 2
electrons only, and no two electrons can have the same
quantum numbers.
ms = Spin quantum number = ±1/2
CHEM& 141 F08
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1
Filling in Electrons
The Rules:
• Aufbau Principle: Electrons fill in electrons by order
of energy, from low → high.
– Not all orbitals are available for all energy (n) levels.
• Pauli Exclusion Principle: Each individual orbital
takes 2 electrons only!
–
–
–
–
There is one s orbital = 2 electrons.
There are three p orbitals = 6 electrons.
There are five d orbitals = 10 electrons.
There are seven f orbitals = 14 electrons.
• Hund’s Rule: If there are multiple orbitals at the
same energy, they fill singly first, before electrons
pair.
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Filling in Electrons
n=1
→ 1s orbital → 2 electrons
n=2
→ 2s orbital → 2 electrons
→ 2p orbital → 6 electrons
n=3
→ 3s orbital → 2 electrons
→ 3p orbital → 6 electrons
→ 3d orbital → 10 electrons
n=4
→ 4s orbital → 2 electrons
→ 4p orbital → 6 electrons
→ 4d orbital →10 electrons
→ 4f orbital → 14 electrons
CHEM& 141 F08
Lower energy rows
have fewer orbitals
available, therefore
there are fewer
elements there!
Beyond n = 4, all
levels have s, p, d
and f orbitals.
4
2
Filling in Electrons
n=1
→ 1s orbital → 2 electrons
n=2
→ 2s orbital → 2 electrons
→ 2p orbital → 6 electrons
n=3
→ 3s orbital → 2 electrons
→ 3p orbital → 6 electrons
→ 3d orbital → 10 electrons
n=4
→ 4s orbital → 2 electrons
→ 4p orbital → 6 electrons
→ 4d orbital →10 electrons
→ 4f orbital → 14 electrons
Lower energy rows
have fewer orbitals
available, therefore
there are fewer
elements there!
Beyond n = 4, all
levels have s, p, d
and f orbitals.
CHEM& 141 F08
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Filling in Electrons
Electrons get filled into orbitals individually:
s:
unoccupied
orbital
orbital with
1 electron
orbital with
2 electrons
The Pauli
Exclusion
Principle!
p:
One electron
Two electrons
Three electrons
Four electrons
Five electrons
Six electrons
Hund’s Rule: fill orbitals singly first, then start pairing!
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3
The Aufbau Principle
• As electrons get added to
elements, the get inserted into
the orbitals in order of energy.
This is not in numerical order!
• The diagram at right shows
the order that electrons fill. To
create the diagram:
– List the orbitals in order.
– Then, draw diagonal lines
downward from right to
left.
– Once you complete a
diagonal, loop back
around.
CHEM& 141 F08
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
Why does this happen?
7
Orbital Energies
Recall: En = -
CHEM& 141 F08
Z2
n2
Rh
8
4
Summary of Electron-filling Rules
• The Pauli Exclusion Principle
– Electrons are like spinning magnets, and must have
opposite alignment.
– We “imagine” this as one “↑” and one “↓”.
• The Aufbau Principle
– The orbitals get filled from the lowest energy to the highest
energy, based on incomplete shielding.
• Hund’s Rule
– When multiple orbitals are present, each orbital gets filled
singly at first, and pairing begins.
– The single electrons all have the same alignment.
CHEM& 141 F08
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Writing Electronic Configurations
To determine the electron configuration:
1) Find the number of electrons for the element.
2) Fill the electrons in order of the Aufbau Principle.
3) Use Hund’s Rule and the Pauli Exclusion
Principle for orbital diagrams.
Example: Nitrogen - Element #7 → 7 electrons
Orbital Diagram
1s
2s
2p
2p
2p
• Each individual orbital gets a “box”.
• Electrons are filled into the boxes
until the total is reached.
CHEM& 141 F08
Electronic Configuration
The number of
electrons within
1s2 2s2 2p2 each set of orbitals
The energy,
or “n” level
The orbital
10
5
Electron Configurations
Determine the orbital diagrams and electron configurations for the
following elements.
He
Li
C
F
Mg
P
Ti
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Electron Configurations
Determine the orbital diagrams and electron configurations for the
following elements.
1s2
He
1s
1s22s1
Li
6e-
1s
2s
1s
2s
2p
2p
2p
1s
2s
2p
2p
2p
C
1s22s22p2
1s22s22p5
F
Mg
1s22s22p63s2
1s
2s
2p
2p
2p
3s
P
1s22s22p63s23p3
1s
2s
2p
2p
2p
3s
3p
3p
3p
Ti
1s
CHEM& 141 F08
2s
2p
2p
2p
3s
3p
3p
3p
1s22s22p63s23p64s23d2
4s
3d
3d
3d
3d
3d
12
6
Electron Configurations and the Periodic
Table
Electron configurations can be read off of the periodic table
CHEM& 141 F08
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Electronic Configurations
Use the periodic table to determine the
electronic configurations for the following
elements.
S
16 e- 1s2 2s2 2p6 3s2 3p4
Ca
V
Ge
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7
Electronic Configuration Shorthand
Consider the electronic for Argon and Calcium:
Ar: 1s2 2s2 2p6 3s2 3p6
← As a noble gas, Argon’s orbitals
are completely filled.
Ca: 1s2 2s2 2p6 3s2 3p6 4s2
[Ar] 4s2 ← We can use the “last” noble gas as a
shorthand in electronic configurations!
The “core”
electrons
The “valence” electrons
Not only are shorthand configurations easier to
write, but they identify the valence electrons, which
are the electrons that are available for reaction!
CHEM& 141 F08
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Electronic Configuration Shorthand
Use the shorthand notation to write the electronic
configurations for the following elements. Also, indicate
the number of valence electrons for each.
K
Mn
As
[Ar] 4s23d104p3 → 15 or 5 valence electrons
Pd
In
Cs
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8
Some Unusual Cases
Consider Cr:
Starting e- config:
Actual:
Consider Ag:
Starting e- config:
Actual:
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Magnetism
1. Give the electron configuration for Fe. (shorthand ok)
2. Can you tell if any of the electrons are paired or unpaired?
3. Draw the orbital box diagram for the valence electrons in Fe.
4. How many unpaired electrons are there in Fe?
Paramagnetism = contains unpaired eDiamagnetism = no unpaired eCHEM& 141 F08
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9
Electronic Configurations for Ions
Let’s consider calcium:
[Ar]4s2
• What is the typical charge on a Calcium
ion?
– Are electrons removed or gained for a
cation?
• How many valence electrons does
calcium have?
What conclusion can you make
from these responses?
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Electronic Configurations for Cations
A cation has fewer electrons than the neutral atom.
• These electrons are removed from the highest “n” level first!
Examples:
Al: [Ne] 3s23p1
Al+3:
Sn: [Kr] 5s24d105p2
Sn+2:
Sn+4:
What do you notice about the resulting cation electron
configurations?
CHEM& 141 F08
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10
Electronic Configurations for Anions
An anion has more electrons than the neutral atom.
• These electrons are added to the atom according to the
Aufbau Principle.
Examples:
P: [Ne] 3s23p3
P-3:
Br: [Ar] 4s23d104p5
Br-1:
What do you notice about the resulting anion electron
configurations? How does this conclusion compare with
the cation configurations?
CHEM& 141 F08
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Electronic Configurations for Ions
Write the electronic configurations for the following
elements and their ions:
Mg/Mg+2:
Fe/Fe+2/Fe+3:Fe [Ar]4s23d6
Fe+2 [Ar]3d6
Fe+3 [Ar]3d5
O/O-2:
Which is more stable,
Fe+2 or Fe+3?
Mn/Mn+2/Mn+7:
CHEM& 141 F08
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11
Electronic Configuration Summary
• For neutral atoms, the number of electrons =
number of protons = atomic number.
– Electrons are inserted according to the Pauli
Exclusion Principle, the Aufbau Principle and
Hund’s Rule.
– Configurations using the shorthand rely on the
previous noble gas and the valence electrons.
• For cations, electrons are removed from at
atom, typically to reach a noble gas
configuration, or other “stable” point.
• For anions, electrons are added to an atom to
reach a noble gas configuration.
CHEM& 141 F08
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Periodic Trends
The location of electrons in orbitals, the shapes of these
orbitals, “metastable” configurations and incomplete
shielding all give rise to a set of trends or
generalizations that we can make using the Periodic
Table.
•
•
•
•
Atomic and Ionic Radii/Size [8.6]
Ionization Energy [8.7]
Electron Affinity [8.8]
Electronegativity [9.6]
All of these govern how atoms combine and react to
form molecules, and govern how molecules interact with
each other.
CHEM& 141 F08
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12
Atomic/Ionic Radii
How does O2- compare to O?
How does Ca+2 compare to Ca?
CHEM& 141 F08
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Ionization Energy
The energy required to completely remove an electron
from the valence shell of an atom/ion, in the gas phase.
M + IE → M+ + e-
Why are there
several deviations?
CHEM& 141 F08
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13
Ionization Energies
Z Element First IE (kJ/mol)
3
4
5
Li
Be
B
Second IE (kJ/mol)
520
899
801
7300
1757
2430
Write the electronic configurations for Be and B. Why is it easier to
remove an electron from B?
CHEM& 141 F08
When predicting IE effects, *ALWAYS*
consider electronic configurations!
27
Electron Affinities
Energy released when and electron is gained by an
atom or an ion in the gas phase.
M + e- → M- + EA
CHEM& 141 F08
http://www.chm.davidson.edu/ronutt/che115/ea.gif
28
14
Electronegativity [Sec. 9.6]
A measure of the tendency of an atom to pull electron
density from another atom to which it is bonded.
CHEM& 141 F08
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Periodic Trends - Summary
CHEM& 141 F08
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15
Concept
Which of the following has the higher:
A
B
Electronegativity?
H
C
Atomic Radius?
Cs
Rb
Electron Affinity?
O
F
Ionization Energy?
O
N
Electronegativity?
H
B
CHEM& 141 F08
A
B
Electron Affinity?
Cl
I
Ionization Energy?
Mg
Al
Atomic Radius?
P
As
Electronegativity?
N
P
31
16