Ch. 9 - Electron Organization
• The Bohr Model [9.4]
• Orbitals [9.5, 9.6]
• Counting Electrons, configurations [9.7]
– Predicting ion charges from electron
configurations.
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Organization of Electrons
• Early ideas of electron organization pointed
to the “cloud” - nucleus in the center,
electrons freely floating around.
– If this model were true, an “excited” element
would emit a full spectrum of light.
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Organization of Electrons
• We have observed that when elements are
excited, they emit discrete “lines” of light.
– The emission is caused by electrons moving
about.
– The discrete lines indicate additional organization
of electrons - it’s not just a cloud!
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Organization of Electrons - The Bohr Model
• Developed in 1913 by Niels Bohr
• Electrons lie in discrete, spherical orbitals
(n level) around the central nucleus.
• Postulates:
– Energy-level: An electron
can only have specific
energy values in an atom.
– Transitions: An electron
can change energy only
by going from one energy
level to another.
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Over time, modifications to the Bohr Model led
to the organization of electrons in orbitals.
4
2
The Orbitals
The s orbital
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The Orbitals
The p orbitals
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The Orbitals
The d orbitals
For a cool representation of all of the orbitals from n = 0 thru n = 10, check out: The Grand Orbital Table
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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
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.
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Filling in Electrons
Electrons get filled into orbitals individually:
s:
p:
unoccupied
orbital
orbital with
1 electron
orbital with
2 electrons
The Pauli
Exclusion
Principle!
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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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.
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
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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.
• 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.
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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.
Electronic Configuration
The number of
electrons within
2
2
2
1s 2s 2p each set of orbitals
The energy,
or “n” level
The orbital
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Electron Configurations
Determine the orbital diagrams and electron configurations for the
following elements.
He
Li
C
F
Mg
P
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Electron Configurations
Determine the orbital diagrams and electron configurations for the
following elements.
He
1s2
1s
1s22s1
Li
6e-
1s
2s
1s
2s
C
2p
2p
2p
1s22s22p2
1s22s22p5
F
1s
2s
2p
2p
2p
Mg
1s22s22p63s2
1s
2s
2p
2p
2p
3s
1s
2s
2p
2p
2p
3s
P
1s22s22p63s23p3
3p
3p
3p
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Electron Configurations and the Periodic Table
Electron configurations can be read off of the periodic table
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Electronic Configurations
Use the periodic table to determine the electronic
configurations for the following elements.
Be
S
16 e-
1s2 2s2 2p6 3s2 3p4
Ca
V
Ge
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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!
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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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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?
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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?
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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:
Mn/Mn+2/Mn+7:
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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.
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Periodic Trends
Based on similarities in electronic
configurations within a group and a period,
some generalizations and trends can be
predicted.
 Atomic Size
 Ionization Energy
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Atomic Size
• This trend typically looks at the atomic
radius, but also applies to the volume.
• Each atom is treated as a marble - a hard
sphere.
• The size is loosely based on the valence
shell, or the “n-level”
– The radius increases as you go down ↓ a group.
• The size is determined by the number of
protons available to attract the electrons.
– The radius decreases as you across → a period.
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Atomic Size
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Atomic Size
Choose the larger atom in each pair:
C or O
Li or K
P or Al
Br or I
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Ionization Energy
• Defined as the amount of energy required to
remove an electron from an atom in the gas
phase.
M(g) + IE → M+1(g) + e-(g)
• Electrons are always removed from the
valence shell.
• The IE tends to decrease down ↓ a group.
– Electrons get further away from the nucleus!
• The IE tends to increase across → a period.
– Electrons are closer to the nucleus!
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Ionization Energy
Choose the atom with the lower ionization
energy in each pair:
C or O
Li or K
P or Al
Br or I
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Electron Affinity
• Defined as the amount of energy released
from an atom in the gas phase when an
electron is captured.
M(g) + e- → M-1(g) + EA
• Electrons are always added to the last
unfilled orbital.
• The EA tends to decrease down ↓ a group.
– The electron spends less time near the nucleus.
• The EA tends to increase across → a period.
– The addition of an electron gets the atom closer
to a noble gas.
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Electron Affinity
Choose the atom with the higher in each pair:
N or F
Ca or Mg
P or Al
Br or I
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Periodic Trend Summary
EA increases
IE increases
Radius decreases
Radius increases
IE decreases
EA decreases
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