Download ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY

ATOMIC ELECTRON
CONFIGURATIONS AND
PERIODICITY
1
Arrangement of
Electrons in Atoms
Electrons in atoms are arranged as
SHELLS (n)
SUBSHELLS (l)
ORBITALS (ml)
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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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4
Electron
Spin
Quantum
Number,
ms
Can be proved experimentally that electron
has a spin. Two spin directions are given by
ms where ms = +1/2 and -1/2.
Electron Spin Quantum Number
Diamagnetic: NOT attracted to a magnetic
field
Paramagnetic: substance is attracted to a
magnetic field. Substance has unpaired
electrons.
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QUANTUM
NUMBERS
n ---> shell
1, 2, 3, 4, ...
l ---> subshell
0, 1, 2, ... n - 1
ml ---> orbital
-l ... 0 ... +l
ms ---> 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.
Electrons in Atoms
When n = 1, then l = 0
this shell has a single orbital (1s) to
which 2e- can be assigned.
When n = 2, then l = 0, 1
2s orbital
2e-
three 2p orbitals
6e-
TOTAL =
8e-
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Electrons in Atoms
When n = 3, then l = 0, 1, 2
3s orbital
three 3p orbitals
five 3d orbitals
TOTAL =
2e6e10e18e-
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Electrons in Atoms
When n = 4, then l = 0, 1, 2, 3
4s orbital
2ethree 4p orbitals
6efive 4d orbitals
10eseven 4f orbitals
14eTOTAL =
32e-
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 l.
•
See Figure 8.5, page 295 and Screen 8. 7.
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 n + l
increases.
b) for subshells of same
n + l, subshell with
lower n is lower in
energy.
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Electron
Filling
Order
Figure 8.5
Effective Nuclear Charge, Z*
• Z* is the nuclear charge experienced by
the outermost electrons. See p. 295 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
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
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” for Electron Configuration tool.
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Electron Configurations
and the Periodic Table
Figure 8.7
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
Group 2A
Atomic number = 4
1s22s2 ---> 4 total
electrons
Boron
3p
3s
2p
2s
1s
Group 3A
Atomic number = 5
1s2 2s2 2p1 --->
5 total electrons
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Carbon
3p
3s
2p
2s
1s
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
Group 5A
Atomic number = 7
1s2 2s2 2p3 --->
7 total electrons
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Oxygen
3p
3s
2p
2s
1s
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
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
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
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.
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.
Relationship of Electron
Configuration and Region of
the Periodic Table
•
•
•
•
Gray = s block
Orange = p block
Green = d block
Violet = f block
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Electron Configurations
and the Periodic Table
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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
Iron
Copper
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Transition Element
Configurations
3d orbitals used
for Sc-Zn (Table
8.4)
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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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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- ---> P3+ [Ne] 3s2 3p0
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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- ---> 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 ---> Fe2+ [Ar] 4s0 3d6
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Ion Configurations
For transition metals, remove ns electrons and
then (n - 1) electrons.
Fe [Ar] 4s2 3d6
loses 2 electrons ---> Fe2+ [Ar] 4s0 3d6
Fe2+
Fe
4s
3d
4s
3d
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Ion Configurations
For transition metals, remove ns electrons and
then (n - 1) electrons.
Fe [Ar] 4s2 3d6
loses 2 electrons ---> Fe2+ [Ar] 4s0 3d6
Fe2+
Fe
4s
3d
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.
Ions with UNPAIRED ELECTRONS are
PARAMAGNETIC.
Without unpaired electrons DIAMAGNETIC.
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PERIODIC
TRENDS
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
Figure 8.6
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.
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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
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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 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.
Atomic Radii
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Figure 8.9
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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
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Trends in Atomic Size
See Figures 8.9 & 8.10
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
30
35
40
Sizes of Transition Elements
See Figure 8.10
• 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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Ion Sizes
Li,152 pm
3e and 3p
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.
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.
Trends in Ion Sizes
Figure 8.13
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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?
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-
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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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General Periodic Trends
• Atomic and ionic size
• Ionization energy
• Electron affinity
Higher Z*.
Electrons held
more tightly.
Larger orbitals.
Electrons held less
tightly.
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Atomic Radii
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Figure 8.9
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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
5
7
9
11
Na
13
15
17
19
K
21
23
25
27
29
31
Atomic Number
33
35
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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Periodic Trend in
the Reactivity of
Alkali Metals
with Water
Lithium
Sodium
Potassium
Ionization Energy
See Screen 8.12
Mg (g) + 735 kJ ---> Mg+ (g) + e-
Mg+ (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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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
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Electron Affinity of Oxygen
O atom [He] 
 

+ electron
O- ion [He] 
 
EA = - 141 kJ

∆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

∆E is zero for Ndue to electronelectron
repulsions.
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Trends in Electron Affinity
• See Figure 8.12 and
Appendix F
• Affinity for electron
increases across a
period (EA becomes
more positive).
• Affinity decreases down
a group (EA becomes
less positive).
Atom EA
F
+328 kJ
Cl +349 kJ
Br +325 kJ
I
+295 kJ
Trends in Electron Affinity
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