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John C. Kotz
Paul M. Treichel
John Townsend
http://academic.cengage.com/kotz
Chapter 7
Atomic Electron Configurations
and Chemical Periodicity
John C. Kotz • State University of New York, College at Oneonta
2
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© 2009 Brooks/Cole - Cengage
ATOMIC ELECTRON
CONFIGURATIONS AND
PERIODICITY
© 2009 Brooks/Cole - Cengage
3
Arrangement of
Electrons in Atoms
Electrons in atoms are arranged as
SHELLS (n)
SUBSHELLS (s)
ORBITALS (ms)
© 2009 Brooks/Cole - Cengage
4
Arrangement of
Electrons in Atoms
Each orbital can be assigned no
more than 2 electrons!
This is tied to the existence of a 4th
quantum number, the electron
spin quantum number, ms.
© 2009 Brooks/Cole - Cengage
5
6
PLAY MOVIE
Electron
Spin
Quantum
Number,
ms
Can be proved experimentally that electron
has an intrinsic property referred to as
“spin.” Two spin directions are given by
ms where ms = +1/2 and -1/2.
© 2009 Brooks/Cole - Cengage
Electron Spin and Magnetism
•Diamagnetic: NOT
attracted to a magnetic
field
•Paramagnetic:
PLAY MOVIE
© 2009 Brooks/Cole - Cengage
substance is attracted to
a magnetic field.
•Substances with
unpaired electrons are
paramagnetic.
7
Measuring Paramagnetism
Paramagnetic: substance is attracted to a
magnetic field. Substance has unpaired electrons.
Diamagnetic: NOT attracted to a magnetic field
See Active Figure 6.18
© 2009 Brooks/Cole - Cengage
8
9
QUANTUM NUMBERS
Now there are four!
n f shell
1, 2, 3, 4, ...
s f subshell
0, 1, 2, ... n - 1
ms f orbital
- s ... 0 ... + s
ms f electron spin
+1/2 and -1/2
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10
Pauli Exclusion Principle
No two electrons in the
same atom can have
the same set of 4
quantum numbers.
That is, each electron has a
unique address.
© 2009 Brooks/Cole - Cengage
Electrons in Atoms
When n = 1, then s = 0
this shell has a single orbital (1s) to which 2ecan be assigned.
When n = 2, then s = 0, 1
2s orbital
2e-
three 2p orbitals
6e-
TOTAL =
8e-
© 2009 Brooks/Cole - Cengage
11
Electrons in Atoms
When n = 3, then s = 0, 1, 2
3s orbital
2ethree 3p orbitals
6efive 3d orbitals
10eTOTAL =
18e-
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12
Electrons in Atoms
When n = 4, then s = 0, 1, 2, 3
4s orbital
three 4p orbitals
five 4d orbitals
seven 4f orbitals
TOTAL =
© 2009 Brooks/Cole - Cengage
2e6e10e14e32e-
And many more!
13
14
© 2009 Brooks/Cole - Cengage
15
Assigning Electrons to Atoms
• Electrons generally assigned to orbitals of
successively higher energy.
• For H atoms, E = - C(1/n2). E depends only
on n.
• For many-electron atoms, energy depends
on both n and s.
•
See Active Figure 7.1 and Figure 7.2
© 2009 Brooks/Cole - Cengage
Assigning Electrons to Subshells
PLAY MOVIE
© 2009 Brooks/Cole - Cengage
• In H atom all subshells
of same n have same
energy.
• In many-electron atom:
a) subshells increase in
energy as value of n + s
increases.
b) for subshells of same n
+ s, subshell with lower
n is lower in energy.
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17
Electron
Filling
Order
See Figure 7.2
© 2009 Brooks/Cole - Cengage
Effective Nuclear Charge, Z*
• Z* is the nuclear charge experienced by
the outermost electrons. See Figure 7.3
• Explains why E(2s) < E(2p)
• Z* increases across a period owing to
incomplete shielding by inner electrons.
• Estimate Z* = [ Z - (no. inner electrons) ]
• Charge felt by 2s e- in Li
Z* = 3 - 2 = 1
• Be
Z* = 4 - 2 = 2
• B
Z* = 5 - 2 = 3
and so on!
© 2009 Brooks/Cole - Cengage
18
19
Effective
Nuclear
Charge
See Figure 7.3
Z* is the nuclear
charge experienced
by the outermost
electrons.
Electron cloud
for 1s electrons
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20
Writing Atomic Electron
Configurations
Two ways of
writing configs.
One is called
the spdf
notation.
spdf notation
for H, atomic number = 1
1
1s
value of n
© 2009 Brooks/Cole - Cengage
no. of
electrons
value of l
Writing Atomic Electron
Configurations
Two ways of
writing
configs. Other
is called the
orbital box
notation.
ORBITAL BOX NOTATION
for He, atomic number = 2
Arrows
2
depict
electron
spin
1s
1s
One electron has n = 1, s = 0, ms = 0, ms = + 1/2
Other electron has n = 1, s = 0, ms = 0, ms = - 1/2
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21
22
See “Toolbox” in ChemNow for Electron Configuration tool.
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Electron Configurations
and the Periodic Table
23
See Active Figure 7.4
© 2009 Brooks/Cole - Cengage
Lithium
Group 1A
Atomic number = 3
1s22s1 f 3 total electrons
3p
3s
2p
2s
1s
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25
Beryllium
3p
3s
2p
2s
1s
© 2009 Brooks/Cole - Cengage
Group 2A
Atomic number = 4
1s22s2 f 4 total
electrons
Boron
3p
3s
2p
2s
1s
© 2009 Brooks/Cole - Cengage
Group 3A
Atomic number = 5
1s2 2s2 2p1 f
5 total electrons
26
Carbon
Group 4A
Atomic number = 6
1s2 2s2 2p2 f
6 total electrons
3p
3s
2p
2s
1s
© 2009 Brooks/Cole - Cengage
Here we see for the first time
HUND’S RULE. When
placing electrons in a set of
orbitals having the same
energy, we place them singly
as long as possible.
27
28
Nitrogen
3p
3s
2p
2s
1s
© 2009 Brooks/Cole - Cengage
Group 5A
Atomic number = 7
1s2 2s2 2p3 f
7 total electrons
29
Oxygen
3p
3s
2p
2s
1s
© 2009 Brooks/Cole - Cengage
Group 6A
Atomic number = 8
1s2 2s2 2p4 f
8 total electrons
30
Fluorine
3p
3s
2p
2s
1s
© 2009 Brooks/Cole - Cengage
Group 7A
Atomic number = 9
1s2 2s2 2p5 f
9 total electrons
31
Neon
Group 8A
Atomic number = 10
1s2 2s2 2p6 f
10 total electrons
3p
3s
2p
2s
1s
© 2009 Brooks/Cole - Cengage
Note that we have
reached the end of
the 2nd period, and
the 2nd shell is full!
32
Electron Configurations of
p-Block Elements
PLAY MOVIE
© 2009 Brooks/Cole - Cengage
Sodium
Group 1A
Atomic number = 11
1s2 2s2 2p6 3s1 or
“neon core” + 3s1
[Ne] 3s1 (uses rare gas notation)
Note that we have begun a new period.
All Group 1A elements have
[core]ns1 configurations.
© 2009 Brooks/Cole - Cengage
33
34
Aluminum
Group 3A
Atomic number = 13
1s2 2s2 2p6 3s2 3p1
[Ne] 3s2 3p1
All Group 3A elements
have [core] ns2 np1
configurations where n
is the period number.
3p
3s
2p
2s
1s
© 2009 Brooks/Cole - Cengage
35
Phosphorus
Yellow P
Group 5A
Atomic number = 15
1s2 2s2 2p6 3s2 3p3
[Ne] 3s2 3p3
All Group 5A elements
have [core ] ns2 np3
configurations where n
is the period number.
Red P
3p
3s
2p
2s
1s
© 2009 Brooks/Cole - Cengage
36
Calcium
Group 2A
Atomic number = 20
1s2 2s2 2p6 3s2 3p6 4s2
[Ar] 4s2
All Group 2A elements have
[core]ns2 configurations where n
is the period number.
© 2009 Brooks/Cole - Cengage
Electron Configurations
and the Periodic Table
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37
Transition Metals
Table 7.4
All 4th period elements have the
configuration [argon] nsx (n - 1)dy
and so are d-block elements.
Chromium
© 2009 Brooks/Cole - Cengage
Iron
Copper
38
Transition Element
Configurations
3d orbitals used
for Sc-Zn (Table
7.4)
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39
40
© 2009 Brooks/Cole - Cengage
Lanthanides and Actinides
All these elements have the configuration
[core] nsx (n - 1)dy (n - 2)fz and so are
f-block elements.
Cerium
[Xe] 6s2 5d1 4f1
Uranium
[Rn] 7s2 6d1 5f3
© 2009 Brooks/Cole - Cengage
41
Lanthanide Element
Configurations
4f orbitals used for
Ce - Lu and 5f for
Th - Lr (Table 7.2)
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42
43
© 2009 Brooks/Cole - Cengage
Ion Configurations
To form cations from elements remove 1 or
more e- from subshell of highest n [or
highest (n + l)].
P [Ne] 3s2 3p3 - 3e- f P3+ [Ne] 3s2 3p0
3p
3p
3s
3s
2p
2p
2s
2s
1s
1s
© 2009 Brooks/Cole - Cengage
44
Ion Configurations
For transition metals, remove ns electrons and
then (n - 1) electrons.
Fe [Ar] 4s2 3d6
loses 2 electrons f Fe2+ [Ar] 4s0 3d6
Fe2+
Fe
4s
3d
To form cations, always
remove electrons of
highest n value first!
© 2009 Brooks/Cole - Cengage
4s
3d
Fe3+
4s
3d
45
Ion Configurations
How do we know the configurations of ions?
Determine the magnetic properties of ions.
Sample
of Fe2O3
© 2009 Brooks/Cole - Cengage
Sample
of Fe2O3
with
strong
magnet
46
Ion Configurations
How do we know the configurations of ions?
Determine the magnetic properties of ions.
Ions with UNPAIRED ELECTRONS are
PARAMAGNETIC.
Without unpaired electrons DIAMAGNETIC.
Fe3+ ions in Fe2O3
have 5 unpaired
electrons and make
the sample
paramagnetic.
© 2009 Brooks/Cole - Cengage
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48
PLAY MOVIE
PLAY MOVIE
PERIODIC
TRENDS
© 2009 Brooks/Cole - Cengage
PLAY MOVIE
General Periodic Trends
• Atomic and ionic size
• Ionization energy
• Electron affinity
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.
© 2009 Brooks/Cole - Cengage
49
Effective Nuclear Charge, Z*
• Z* is the nuclear charge experienced by
the outermost electrons. See Figure 7.3
• Explains why E(2s) < E(2p)
• Z* increases across a period owing to
incomplete shielding by inner electrons.
• Estimate Z* = [ Z - (no. inner electrons) ]
• Charge felt by 2s e- in Li
Z* = 3 - 2 = 1
• Be
Z* = 4 - 2 = 2
• B
Z* = 5 - 2 = 3
and so on!
© 2009 Brooks/Cole - Cengage
50
51
Effective
Nuclear
Charge
See Figure 7.3
Z* is the nuclear
charge experienced
by the outermost
electrons.
Electron cloud
for 1s electrons
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52
Effective Nuclear Charge Z*
The 2s electron PENETRATES the region
occupied by the 1s electron.
2s electron experiences a higher positive
charge than expected.
PLAY MOVIE
© 2009 Brooks/Cole - Cengage
Effective Nuclear Charge, Z*
• Atom
•
•
•
•
•
•
•
Li
Be
B
C
N
O
F
Z* Experienced by Electrons in
Valence Orbitals
+1.28
------+2.58
Increase in
+3.22
Z* across a
+3.85
period
+4.49
+5.13
[Values calculated using Slater’s Rules]
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53
Orbital Energies
Orbital energies “drop” as Z* increases
ChemNow Screens 8.9 - 8.13, Simulations
© 2009 Brooks/Cole - Cengage
54
General Periodic Trends
• Atomic and ionic size
• Ionization energy
• Electron affinity
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.
© 2009 Brooks/Cole - Cengage
55
56
Atomic Radii
See Active Figure 7.8
© 2009 Brooks/Cole - Cengage
57
Atomic Size
• Size goes UP on going down
a group. See Figure 7.8.
• Because electrons are
added further from the
nucleus, there is less
attraction.
• Size goes DOWN on going
across a period.
© 2009 Brooks/Cole - Cengage
58
Atomic Size
Size decreases across a period owing
to increase in Z*. Each added electron
feels a greater and greater + charge.
Large
Small
Increase in Z*
© 2009 Brooks/Cole - Cengage
59
Trends in Atomic Size
See Active Figure 7.8
Radius (pm)
250
K
1st transition
series
3rd period
200
Na
2nd period
Li
150
Kr
100
Ar
Ne
50
He
0
0
5
10
15
20
25
Atomic Number
© 2009 Brooks/Cole - Cengage
30
35
40
Sizes of Transition Elements
See Figure 7.9
© 2009 Brooks/Cole - Cengage
60
Sizes of Transition Elements
See Figure 7.9
• 3d subshell is inside the 4s
subshell.
• 4s electrons feel a more or less
constant Z*.
• Sizes stay about the same and
chemistries are similar!
© 2009 Brooks/Cole - Cengage
61
Density of Transition Metals
25
20
6th period
Density (g/mL)
15
10
5th period
4th period
5
0
3B
4B
5B
6B
7B
8B
Group
© 2009 Brooks/Cole - Cengage
1B
2B
62
Ion Sizes
Li,152 pm
3e and 3p
© 2009 Brooks/Cole - Cengage
Does+ the size go
up+ or down
Li , 60 pm
when
an
2e and 3losing
p
electron to form
a cation?
63
64
Ion Sizes
+
Li,152 pm
3e and 3p
Li + , 78 pm
2e and 3 p
Forming
a cation.
• CATIONS are SMALLER than the
atoms from which they come.
• The electron/proton attraction has
gone UP and so size DECREASES.
© 2009 Brooks/Cole - Cengage
Ion Sizes
Does the size go up or
down when gaining an
electron to form an
anion?
© 2009 Brooks/Cole - Cengage
65
66
Ion Sizes
F, 71 pm
9e and 9p
F- , 133 pm
10 e and 9 p
Forming
an anion.
• ANIONS are LARGER than the atoms from
which they come.
• The electron/proton attraction has gone
DOWN and so size INCREASES.
• Trends in ion sizes are the same as atom
sizes.
© 2009 Brooks/Cole - Cengage
Trends in Ion Sizes
67
See Active Figure 7.12
© 2009 Brooks/Cole - Cengage
68
Redox Reactions
Why do metals lose
electrons in their
reactions?
Why does Mg form Mg2+
ions and not Mg3+?
Why do nonmetals take
on electrons?
© 2009 Brooks/Cole - Cengage
Ionization Energy
IE = energy required to remove an electron
from an atom in the gas phase.
PLAY MOVIE
Mg (g) + 738 kJ f Mg+ (g) + e-
© 2009 Brooks/Cole - Cengage
69
Ionization Energy
IE = energy required to remove an electron
from an atom in the gas phase.
Mg (g) + 738 kJ f Mg+ (g) + e-
PLAY MOVIE
Mg+ (g) + 1451 kJ f Mg2+ (g) + eMg+ has 12 protons and only 11
electrons. Therefore, IE for Mg+ > Mg.
© 2009 Brooks/Cole - Cengage
70
Ionization Energy
Mg (g) + 735 kJ f Mg+ (g) + eMg+ (g) + 1451 kJ f Mg2+ (g) + e-
PLAY MOVIE
Mg2+ (g) + 7733 kJ f Mg3+ (g) + eEnergy cost is very high to dip into a
shell of lower n.
This is why ox. no. = Group no.
© 2009 Brooks/Cole - Cengage
71
Trends in Ionization Energy
See Active Figure 7.10
© 2009 Brooks/Cole - Cengage
73
74
Trends in Ionization Energy
1st Ionization energy (kJ/mol)
2500
He
Ne
2000
Ar
1500
Kr
1000
500
0
1
H
3
Li
© 2009 Brooks/Cole - Cengage
5
7
9
11
Na
13
15
17
19
K
21
23
25
27
29
31
Atomic Number
33
35
Orbital Energies
As Z* increases, orbital energies
“drop” and IE increases.
CD-ROM Screens 8.9 - 8.13, Simulations
© 2009 Brooks/Cole - Cengage
75
Trends in Ionization Energy
• IE increases across a period
because Z* increases.
• Metals lose electrons more
easily than nonmetals.
• Metals are good reducing
agents.
• Nonmetals lose electrons with
difficulty.
© 2009 Brooks/Cole - Cengage
76
77
Trends in Ionization Energy
• IE decreases down a
group
• Because size increases.
• Reducing ability
generally increases down
the periodic table.
• See reactions of Li, Na, K
© 2009 Brooks/Cole - Cengage
78
PLAY MOVIE
Periodic Trend in
the Reactivity of
Alkali Metals
with Water
Lithium
PLAY MOVIE
Sodium
© 2009 Brooks/Cole - Cengage
PLAY MOVIE
Potassium
79
Electron Affinity
A few elements GAIN electrons to
form anions.
Electron affinity is the energy
involved when an atom gains
an electron to form an anion.
A(g) + e- f A-(g) E.A. = ∆U
© 2009 Brooks/Cole - Cengage
80
Electron Affinity of Oxygen
O atom [He] 
 

+ electron
O- ion [He] 
 
EA = - 141 kJ
© 2009 Brooks/Cole - Cengage

∆U is EXOthermic
because O has
an affinity for an
e-.
81
Electron Affinity of Nitrogen
N atom [He] 
 

+ electron
N- ion
[He] 


EA = 0 kJ
© 2009 Brooks/Cole - Cengage

∆U is zero for Ndue to electronelectron
repulsions.
Trends in Electron Affinity
See Active Figure 7.11
© 2009 Brooks/Cole - Cengage
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83
Trends in Electron Affinity
• See Figure 7.11 and
Appendix F
• Affinity for electron
increases across a
period (EA becomes
more positive).
• Affinity decreases down
a group (EA becomes
less positive).
© 2009 Brooks/Cole - Cengage
Atom EA
F
+328 kJ
Cl +349 kJ
Br +325 kJ
I
+295 kJ
Note effect of atom
size on F vs. Cl