Transcript Chapter 7

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
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ATOMIC ELECTRON
CONFIGURATIONS AND
PERIODICITY
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Arrangement of
Electrons in Atoms
Electrons in atoms are arranged as
SHELLS (n)
SUBSHELLS (s)
ORBITALS (ms)
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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.
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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.
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Electron Spin and Magnetism
•Diamagnetic: NOT
attracted to a magnetic
field
•Paramagnetic:
PLAY MOVIE
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substance is attracted to
a magnetic field.
•Substances with
unpaired electrons are
paramagnetic.
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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
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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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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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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-
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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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Electrons in Atoms
When n = 4, then s = 0, 1, 2, 3
4s orbital
three 4p orbitals
five 4d orbitals
seven 4f orbitals
TOTAL =
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2e6e10e14e32e-
And many more!
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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
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Assigning Electrons to Subshells
PLAY MOVIE
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• 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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Electron
Filling
Order
See Figure 7.2
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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!
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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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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, s = 0, ms = 0, ms = + 1/2
Other electron has n = 1, s = 0, ms = 0, ms = - 1/2
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See “Toolbox” in ChemNow for Electron Configuration tool.
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Electron Configurations
and the Periodic Table
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See Active Figure 7.4
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Lithium
Group 1A
Atomic number = 3
1s22s1 f 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 f 4 total
electrons
Boron
3p
3s
2p
2s
1s
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Group 3A
Atomic number = 5
1s2 2s2 2p1 f
5 total electrons
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Carbon
Group 4A
Atomic number = 6
1s2 2s2 2p2 f
6 total electrons
3p
3s
2p
2s
1s
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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 f
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 f
8 total electrons
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Fluorine
3p
3s
2p
2s
1s
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Group 7A
Atomic number = 9
1s2 2s2 2p5 f
9 total electrons
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Neon
Group 8A
Atomic number = 10
1s2 2s2 2p6 f
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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Electron Configurations of
p-Block Elements
PLAY MOVIE
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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.
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Electron Configurations
and the Periodic Table
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Transition Metals
Table 7.4
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
7.4)
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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 7.2)
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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
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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!
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4s
3d
Fe3+
4s
3d
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Ion Configurations
How do we know the configurations of ions?
Determine the magnetic properties of ions.
Sample
of Fe2O3
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Sample
of Fe2O3
with
strong
magnet
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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.
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PLAY MOVIE
PLAY MOVIE
PERIODIC
TRENDS
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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.
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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!
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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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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
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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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Orbital Energies
Orbital energies “drop” as Z* increases
ChemNow Screens 8.9 - 8.13, Simulations
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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
See Active Figure 7.8
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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.
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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 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
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Sizes of Transition Elements
See Figure 7.9
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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!
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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
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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.
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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.
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Trends in Ion Sizes
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See Active Figure 7.12
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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?
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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-
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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.
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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.
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Trends in Ionization Energy
See Active Figure 7.10
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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
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5
7
9
11
Na
13
15
17
19
K
21
23
25
27
29
31
Atomic Number
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Orbital Energies
As Z* increases, orbital energies
“drop” and IE increases.
CD-ROM Screens 8.9 - 8.13, Simulations
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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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PLAY MOVIE
Periodic Trend in
the Reactivity of
Alkali Metals
with Water
Lithium
PLAY MOVIE
Sodium
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PLAY MOVIE
Potassium
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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- f A-(g) E.A. = ∆U
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Electron Affinity of Oxygen
O atom [He] 
 

+ electron
O- ion [He] 
 
EA = - 141 kJ
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
∆U 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
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
∆U is zero for Ndue to electronelectron
repulsions.
Trends in Electron Affinity
See Active Figure 7.11
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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).
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Atom EA
F
+328 kJ
Cl +349 kJ
Br +325 kJ
I
+295 kJ
Note effect of atom
size on F vs. Cl