Transcript atomic electron configurations and periodicity

Chemistry and Chemical Reactivity
6th Edition
1
John C. Kotz
Paul M. Treichel
Gabriela C. Weaver
CHAPTER 8
Atomic Electron Configurations
and Chemical Periodicity
Lectures written by John Kotz
©2006
2006
Brooks/Cole
Thomson
©
Brooks/Cole
- Thomson
ATOMIC ELECTRON
CONFIGURATIONS AND
PERIODICITY
© 2006 Brooks/Cole - Thomson
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Arrangement of
Electrons in Atoms
Electrons in atoms are arranged as
SHELLS (n)
SUBSHELLS (l)
ORBITALS (ml)
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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.
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Assigning Electrons to Subshells
• In H atom all subshells
of same n have same
energy.
• In many-electron atom:
a) subshells increase in
energy as value of 
increases. S<p<d<f
b) for shells increase in
energy: 1<2<3<4<…
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Electron
Filling
Order
Figure 8.5
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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
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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, l = 0, ml = 0, ms = + 1/2
Other electron has n = 1, l = 0, ml = 0, ms = - 1/2
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See “Toolbox” on CD for Electron Configuration tool.
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Electron Configurations
and the Periodic Table
Active Figure 8.7
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The electron configuration
for chlorine is:
1.
2.
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3
4
5
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4.
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3.4262
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0% 0% 0% 0%
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Lithium
Group 1A
Atomic number = 3
1s22s1 ---> 3 total electrons
3p
3s
2p
2s
1s
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Beryllium
3p
3s
2p
2s
1s
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Group 2A
Atomic number = 4
1s22s2 ---> 4 total
electrons
Boron
3p
3s
2p
2s
1s
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Group 3A
Atomic number = 5
1s2 2s2 2p1 --->
5 total electrons
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Carbon
3p
3s
2p
2s
1s
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Group 4A
Atomic number = 6
1s2 2s2 2p2 --->
6 total electrons
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.
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Nitrogen
3p
3s
2p
2s
1s
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Group 5A
Atomic number = 7
1s2 2s2 2p3 --->
7 total electrons
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Oxygen
3p
3s
2p
2s
1s
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Group 6A
Atomic number = 8
1s2 2s2 2p4 --->
8 total electrons
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Fluorine
Group 7A
Atomic number = 9
1s2 2s2 2p5 --->
9 total electrons
3p
3s
2p
2s
1s
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Neon
Group 8A
Atomic number = 10
1s2 2s2 2p6 --->
10 total electrons
3p
3s
2p
2s
1s
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Note that we have
reached the end of
the 2nd period, and
the 2nd shell is full!
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1.
2.
3.
4.
W
Mo
Ru
Pm
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Electron Configurations of
p-Block Elements
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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.
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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
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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
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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.
© 2006 Brooks/Cole - Thomson
Electron Configurations
and the Periodic Table
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1.
2.
3.
4.
Hf
Lu
Pb
Sn
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3
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Transition Metals
Table 8.4
All 4th period elements have the
configuration [argon] nsx (n - 1)dy
and so are d-block elements.
Chromium
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Iron
Copper
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Transition Element
Configurations
3d orbitals used
for Sc-Zn (Table
8.4)
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1.
2.
3.
4.
Zn
Ca
Ge
Ni
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The electron configuration
for tin, Sn, element 50, is:
1.
2.
3.
4.
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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
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Lanthanide Element
Configurations
4f orbitals used for
Ce - Lu and 5f for
Th - Lr (Table 8.2)
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PERIODIC
TRENDS
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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.
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Effective Nuclear Charge, Z*
• Z* is the nuclear charge experienced by the
outermost electrons. See Figure 8.6 and and Screen 8.6.
• Explains why E(2s) < E(2p)
• Z* increases across a period owing to incomplete
shielding by inner electrons.
• Estimate Z* by --> [ 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!
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Effective
Nuclear
Charge
Figure 8.6
Z* is the nuclear
charge experienced
by the outermost
electrons.
Electron cloud
for 1s electrons
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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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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.
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Atomic Radii
Active Figure 8.11
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Atomic Size
• Size goes UP on going down
a group. See Figure 8.9.
• Because electrons are
added further from the
nucleus, there is less
attraction.
• Size goes DOWN on going
across a period.
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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*
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Trends in Atomic Size
See Active Figure 8.11
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
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35
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Sizes of Transition Elements
See Figure 8.12
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Sizes of Transition Elements
See Figure 8.12
• 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!
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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
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1B
2B
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Ion Sizes
Li,152 pm
3e and 3p
© 2006 Brooks/Cole - Thomson
Does+ the size go
up+ or down
Li , 60 pm
when
an
2e and 3losing
p
electron to form
a cation?
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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.
© 2006 Brooks/Cole - Thomson
Ion Sizes
Does the size go up or
down when gaining an
electron to form an
anion?
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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.
© 2006 Brooks/Cole - Thomson
Trends in Ion Sizes
Active Figure 8.15
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Compare the elements Na, B, Al, and C with regard
to the following properties: Which has the largest
atomic radius?
1.
2.
3.
4.
Na
B
Al
C
0%
1
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0%
0%
2
3
0%
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Compare the elements Na, B, Al, and C with regard to the
following properties: Which has the largest (most negative)
electron affinity?
1.
2.
3.
4.
Na
B
Al
C
0%
1
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0%
0%
2
3
0%
4
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Compare the elements Na, B, Al, and C with regard to the
following properties: Which has the largest (most positive)
ionization energy?
1.
2.
3.
4.
Na
B
Al
C
0%
1
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0%
0%
2
3
0%
4
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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?
© 2006 Brooks/Cole - Thomson
Ionization Energy
See CD Screen 8.12
IE = energy required to remove an electron
from an atom in the gas phase.
Mg (g) + 738 kJ ---> Mg+ (g) + e-
© 2006 Brooks/Cole - Thomson
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Ionization Energy
See Screen 8.12
IE = energy required to remove an electron
from an atom in the gas phase.
Mg (g) + 738 kJ ---> Mg+ (g) + e-
Mg+ (g) + 1451 kJ ---> Mg2+ (g) + eMg+ has 12 protons and only 11
electrons. Therefore, IE for Mg+ > Mg.
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Ionization Energy
See Screen 8.12
Mg (g) + 735 kJ ---> Mg+ (g) + eMg+ (g) + 1451 kJ ---> Mg2+ (g) + e-
Mg2+ (g) + 7733 kJ ---> Mg3+ (g) + eEnergy cost is very high to dip into a
shell of lower n.
This is why ox. no. = Group no.
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Trends in Ionization Energy
Active Figure 8.13
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Trends in Ionization Energy
1st Ionization energy (kJ/mol)
2500
He
Ne
2000
Ar
1500
Kr
1000
500
0
1
H
3
Li
© 2006 Brooks/Cole - Thomson
5
7
9
11
Na
13
15
17
19
K
21
23
25
27
29
31
Atomic Number
33
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Orbital Energies
As Z* increases, orbital energies
“drop” and IE increases.
CD-ROM Screens 8.9 - 8.13, Simulations
© 2006 Brooks/Cole - Thomson
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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.
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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
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Which of the following groups of elements is arranged
correctly in order of increasing first ionization energy?
1.
2.
3.
4.
Mg < C < N < F
N < Mg < C < F
Mg < N < C < F
F < C < Mg < N
0%
1
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0%
0%
2
3
0%
4
Which of the following elements would have the
greatest difference between the first and the second
ionization energy?
1.
2.
3.
4.
lithium
carbon
fluorine
nitrogen
0%
1
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0%
0%
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3
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Which of the following groups of elements is arranged
correctly in order of increasing first ionization energy?
1.
2.
3.
4.
B < O < Al < F
B < O < F < Al
Al < B < O < F
F < O < B < Al
0%
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Which of the following groups of elements is arranged correctly in
order of increasing affinity for electrons (that is, electron affinity
becomes more negative)?
1.
2.
3.
4.
Mg < S < Al < Cl
Mg < Al < S < Cl
Al < Mg < S < Cl
Cl < S < Mg < Al
0%
1
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Which of the following groups of
elements is arranged correctly in
order of decreasing atomic radius?
1.
2.
3.
4.
Mg > S > Al > Cl
Mg > Al > S > Cl
Al > Mg > S > Cl
Cl > S > Mg > Al
0%
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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- ---> A-(g) E.A. = ∆E
© 2006 Brooks/Cole - Thomson
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Electron Affinity of Oxygen
O atom [He] 
 

+ electron
O- ion [He] 
 
EA = - 141 kJ
© 2006 Brooks/Cole - Thomson

∆E is EXOthermic
because O has
an affinity for an
e-.
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Electron Affinity of Nitrogen
N atom [He] 
 

+ electron
N- ion
[He] 


EA = 0 kJ
© 2006 Brooks/Cole - Thomson

∆E is zero for Ndue to electronelectron
repulsions.
Trends in Electron Affinity
Active Figure 8.14
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Trends in Electron Affinity
• See Figure 8.14 and
Appendix F
• Affinity for electron
increases across a
period (EA becomes
more positive).
• Affinity decreases down
a group (EA becomes
less positive).
© 2006 Brooks/Cole - Thomson
Atom EA
F
+328 kJ
Cl +349 kJ
Br +325 kJ
I
+295 kJ
Note effect of atom
size on F vs. Cl
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Which of the following groups of elements is arranged
correctly in order of increasing first ionization energy?
1.
2.
3.
4.
Mg < C < N < F
N < Mg < C < F
Mg < N < C < F
F < C < Mg < N
0%
1
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0%
0%
2
3
0%
4
Which of the following elements would have the
greatest difference between the first and the second
ionization energy?
1.
2.
3.
4.
lithium
carbon
fluorine
nitrogen
0%
1
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0%
0%
2
3
0%
4
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98
Which of the following groups of elements is arranged
correctly in order of increasing first ionization energy?
1.
2.
3.
4.
B < O < Al < F
B < O < F < Al
Al < B < O < F
F < O < B < Al
0%
1
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0%
0%
2
3
0%
4
99
Which of the following groups of elements is arranged correctly in
order of increasing affinity for electrons (that is, electron affinity
becomes more negative)?
1.
2.
3.
4.
Mg < S < Al < Cl
Mg < Al < S < Cl
Al < Mg < S < Cl
Cl < S < Mg < Al
0%
1
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0%
0%
2
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100
Which of the following groups of
elements is arranged correctly in
order of decreasing atomic radius?
1.
2.
3.
4.
Mg > S > Al > Cl
Mg > Al > S > Cl
Al > Mg > S > Cl
Cl > S > Mg > Al
0%
1
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2
3
0%
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