# The Periodic Table!

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```					                        The Periodic Table!

• 8.3-5: Electron Configurations
• 8.6-8: Periodic Trends [also 9.6]
• Ion configurations and trends are

The Periodic Table holds a wealth of information
about the organization and and development of
properties, now we unlock them!

CHEM& 141 F08                                                    1

Electron Configurations

n=          1       2                       3

l=          0   0       1       0       1             2

ml = 0          0   -1 0 +1     0   -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                                                    2

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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.

CHEM& 141 F08                                                                   3

Filling in Electrons
n=1           → 1s orbital → 2 electrons
Lower energy rows
n=2           → 2s orbital → 2 electrons
have fewer orbitals
→ 2p orbital → 6 electrons         available, therefore
there are fewer
n=3           → 3s orbital → 2 electrons         elements there!
→ 3p orbital → 6 electrons
→ 3d orbital → 10 electrons

n=4           → 4s orbital → 2 electrons
→ 4p orbital → 6 electrons         Beyond n = 4, all
→ 4d orbital →10 electrons         levels have s, p, d
→ 4f orbital → 14 electrons        and f orbitals.

CHEM& 141 F08                                                                   4

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Filling in Electrons
n=1       → 1s orbital → 2 electrons
Lower energy rows
n=2       → 2s orbital → 2 electrons
have fewer orbitals
→ 2p orbital → 6 electrons            available, therefore
there are fewer
n=3       → 3s orbital → 2 electrons            elements there!
→ 3p orbital → 6 electrons
→ 3d orbital → 10 electrons

n=4       → 4s orbital → 2 electrons
→ 4p orbital → 6 electrons
Beyond n = 4, all
→ 4d orbital →10 electrons            levels have s, p, d
→ 4f orbital → 14 electrons           and f orbitals.

CHEM& 141 F08                                                                  5

Filling in Electrons
Electrons get filled into orbitals individually:

s:                                                        The Pauli
Exclusion
unoccupied       orbital with    orbital with     Principle!
orbital        1 electron      2 electrons

p:
One electron        Two electrons      Three electrons

Four electrons     Five electrons        Six electrons
Hund’s Rule: fill orbitals singly first, then start pairing!

CHEM& 141 F08                                                                  6

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The Aufbau Principle
• As electrons get added to
elements, the get inserted into
the orbitals in order of energy.    1s
This is not in numerical order!
• The diagram at right shows          2s     2p
the order that electrons fill. To
3s     3p         3d
create the diagram:
– List the orbitals in order.     4s     4p         4d        4f
– Then, draw diagonal lines
downward from right to         5s     5p         5d        5f
left.
6s     6p         6d
– Once you complete a
diagonal, loop back            7s
around.

CHEM& 141 F08
Why does this happen?                              7

Orbital Energies

Z2
Recall: En = -        Rh
n2

CHEM& 141 F08                                                                  8

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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, 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                                                                            9

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            Electronic Configuration
The number of
electrons within

1s     2s     2p   2p   2p
1s2 2s2 2p2 each set of orbitals

• Each individual orbital gets a “box”.                   The orbital
The energy,
• Electrons are filled into the boxes     or “n” level
until the total is reached.

CHEM& 141 F08                                                                           10

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Electron Configurations
Determine the orbital diagrams and electron configurations for the
following elements.

He

Li

C

F

Mg

P

Ti
CHEM& 141 F08                                                                                                        11

Electron Configurations
Determine the orbital diagrams and electron configurations for the
following elements.

He           1s2
1s

Li                      1s22s1
1s    2s
6e-   C                                                       1s22s22p2
1s    2s             2p        2p        2p

F                                                       1s22s22p5
1s    2s             2p        2p        2p

Mg                                                                1s22s22p63s2
1s         2s         2p        2p        2p        3s
P                                                                                  1s22s22p63s23p3
1s         2s         2p        2p        2p        3s     3p    3p   3p

Ti
1s    2s        2p        2p        2p         3s    3p   3p    3p    4s    3d   3d   3d   3d   3d

CHEM& 141 F08                                       1s22s22p63s23p64s23d2                                            12

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Electron Configurations and the Periodic
Table

Electron configurations can be read off of the periodic table
CHEM& 141 F08                                                        13

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

CHEM& 141 F08                                                        14

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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”            The “valence” electrons
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                                                                  15

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

CHEM& 141 F08                                                                  16

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Some Unusual Cases
Consider Cr:
Starting e- config:

Actual:

Consider Ag:
Starting e- config:

Actual:

CHEM& 141 F08                                                              17

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 e-
Diamagnetism = no unpaired e-

CHEM& 141 F08                                                              18

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?
CHEM& 141 F08                                                                 19

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                                                                 20

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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?

CHEM& 141 F08                                                              21

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                  Which is more stable,
Fe+3 [Ar]3d5                  Fe+2 or Fe+3?
O/O-2:

Mn/Mn+2/Mn+7:

CHEM& 141 F08                                                              22

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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.

CHEM& 141 F08                                                        23

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                                                        24

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How does O2- compare to O?

How does Ca+2 compare to Ca?
CHEM& 141 F08                                                      25

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                                                      26

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Ionization Energies

Z Element First IE (kJ/mol)                      Second IE (kJ/mol)
3         Li               520                                 7300
4         Be               899                                 1757
5         B                801                                 2430
Write the electronic configurations for Be and B. Why is it easier to
remove an electron from B?

When predicting IE effects, *ALWAYS*
CHEM& 141 F08                                                                                   27
consider electronic configurations!

Electron Affinities

Energy released when and electron is gained by an
atom or an ion in the gas phase.

M + e- → M- + EA

http://www.chm.davidson.edu/ronutt/che115/ea.gif
CHEM& 141 F08                                                                                   28

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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                                                   29

Periodic Trends - Summary

CHEM& 141 F08                                                   30

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Concept

Which of the following has the higher:

A         B                         A    B

Electronegativity?             H         C    Electron Affinity?   Cl   I

Atomic Radius?                 Cs        Rb   Ionization Energy?   Mg   Al

Electron Affinity?             O         F    Atomic Radius?       P    As

Ionization Energy?             O         N    Electronegativity?   N    P

Electronegativity?             H         B

CHEM& 141 F08                                                                    31

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