Chapter 8: ATOMIC ELECTRON
CONFIGURATIONS AND
PERIODICITY
1
Arrangement of
Electrons in Atoms
Electrons in atoms are arranged as
SHELLS (n)
SUBSHELLS (l)
ORBITALS (ml)
2
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.
3
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.
5
6
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
7
Pauli Exclusion Principle
No two electrons in the
same atom can have the
same set of 4 quantum
numbers.
That is, each electron in an
atom has a unique address
of quantum numbers.
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-
8
Electrons in Atoms
When n = 3, then l = 0, 1, 2
3s orbital
three 3p orbitals
five 3d orbitals
TOTAL =
2e6e10e18e-
9
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!
10
11
12
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), the subshell with
lower n is lower in
energy.
13
14
Electron
Filling
Order
Figure 8.5
Effective Nuclear Charge, Z*
• Z* is the nuclear charge experienced by
the outermost electrons.
• 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!
15
Effective
Nuclear
Charge
Figure 8.6
Electron cloud
for 1s electrons
16
17
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
18
19
See “Toolbox” for Electron Configuration tool.
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
20
General Periodic Trends
• Atomic and ionic size
• Ionization energy
• Electron affinity
Higher effective nuclear charge.
Electrons held more tightly
Smaller orbitals.
Electrons held more
tightly.
21
22
Atomic Size
• Size goes UP on going down
a group.
• Because electrons are
added farther from the
nucleus, there is less
attraction.
• Size goes DOWN on going
across a period.
Atomic Radii
23
Figure 8.9
24
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
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?
25
26
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?
27
28
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
29
30
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-
31
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.
32
33
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.
34
35
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
36
Electron Affinity
A few elements GAIN electrons to
form anions.
Electron affinity is the energy
change when an electron is added:
A(g) + e- ---> A-(g)
E.A. = ∆E
37
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-.
38
Electron Affinity of Nitrogen
N atom [He] 
 

+ electron
N- ion
[He] 


EA = 0 kJ

∆E is zero for Ndue to electronelectron
repulsions.
39
Trends in Electron Affinity
Atom EA
• Affinity for electron
F
-328 kJ
increases across a
period (EA becomes
Cl -349 kJ
more negative).
Br -325 kJ
• Affinity decreases down I
-295 kJ
a group (EA becomes
less negative).
Trends in Electron Affinity
40