Chapter #8 - Electron Configuration and
Chemical Periodicity
8.1 Development of the Periodic Table
8.2 Characteristics of Many-Electron Atoms
8.3 The Quantum-Mechanical Model and the Periodic Table
8.4 Trends in Some Key Periodic Atomic Properties
8.5 The Connection Between Atomic Structure and
Chemical Reactivity
Mendeleev’s Predicted vs Actual Properties
of Element # 32 - Germanium
Property
Atomic Mass
Appearance
Density
Molar volume
Specific heat capacity
Oxide density
Sulfide formula and
solubility
Chloride formula
(boiling point)
Chloride density
Element preparation
Predicted Properties
72
Gray Metal
5.5 g/cm3
13 cm3 /mol
0.31 J/g K
4.7 g/cm3
ES2; insoluble in
H2O; soluble in
aqueous (NH4)S
ECl4
< 100oC
1.9 g/cm3
reduction of K2EF6
with sodium
Actual Properties
72.59
Gray Metal
5.35 g/cm3
13.22 cm3/mol
0.32 J/g K
4.23 g/cm3
GeS2; insoluble in
H2O; soluble in
aqueous (NH4)S
GeCl4
84oC
1.844 g/cm3
Reduction of
K2GeF6 with sodium
Observing the Effect of Electron Spin
Fig. 8.1
Summary of Quantum Numbers of
Electrons in Atoms
Name
Symbol
Permitted Values
Property
Principal
n
Positive integers (1,2,3, etc.) Orbital energy
(size)
Angular
momentum
l
Integers from 0 to n - 1
Magnetic
ml
Integers from -l to 0 to +l
Spin
ms
+ 1/2 or -1/2
Orbital shape (the l
values 0, 1, 2, and 3
correspond to the s,
p, d, and f orbitals)
Orbital orientation
Direction of e- spin
Table 8.2
Quantum Numbers - I
• 1) Principal Quantum Number = n
•
Also called the “energy “ quantum number, indicates
the approximate distance from the nucleus .
•
Denotes the electron energy shells around the atom, and
is derived directly from the Schrodinger equation.
•
The higher the value of “n” , the greater the Energy of
the orbital, and hence the energy of electrons in that
orbital.
• Positive integer values of n = 1 , 2 , 3 , etc.
Quantum Numbers - II
• 2) Azimuthal
• Denotes the different energy sublevels within the
main level “n”
• Also indicates the shape of the orbitals around the
nucleus.
• Positive interger values of L are : 0
• n=1 , L=0
n=2,
n=3,L=0,1,2
( n-1 )
L = 0 and 1
Quantum Numbers - III
• 3) Magnetic Quantum Number - mL
Also called the orbital orientation Quantum #
• denotes the direction or orientation in a magnetic field
- Or it denotes the different magnetic geometriesound
the nucleus - three dimensional space
• values can be positive and negative (-L
0
+L)
• L = 0 , mL = 0
L =1 , mL = -1,0,+1
L = 2 , mL = -2,-1,0,1,2
Quantum Numbers - IV
• 4) Spin Quantum Number - ms - gives the
spin of the electron + or • The values of the spin are either :
+ 1 / 2 or - 1 / 2
• n =1 L = 0
• n=2 L=0
L=1
•
mL = 0
mL = 0
mL = -1
mL = 0
mL = +1
ms = + 1/ 2 and - 1/ 2
ms = + 1/ 2 and - 1/ 2
ms = + 1/ 2 and - 1/ 2
ms = + 1/ 2 and - 1/ 2
ms = + 1/ 2 and - 1/ 2
Spectral Evidence of Energy-Level Splitting in
Many-Electron Atoms
Fig. 8.2
Fig. 8.3
Fig. 8.4
Pauli Exclusion Principle:
Each electron in an atom must have a unique set
of quantum numbers !
Only two electrons can be described by the same
orbital and these two electrons must have
opposite spin.
As a Result of the Pauli
Exclusion Principle:
• Electrons with the same spin keep
apart in space whereas electrons of
opposite spin may occupy the same
region of space.
Quantum Numbers - V
•
•
•
•
•
•
•
•
•
•
n=1
n=2
n=3
n=4
L=0
L=0
L=1
L=0
L=1
L=2
L=0
L=1
L=2
L=3
mL = 0
ms = + 1/ 2 & - 1/ 2
mL = 0
for all orbitals
mL = -1 , 0 , +1
mL = 0
mL = -1 , 0 , +1
mL = - 2 , -1 , 0 , +1 , +2
mL = 0
mL = -1 , 0 +1
mL = - 2 , -1 , 0 , +1 , +2
mL = - 3 , - 2 , - 1 , 0, +1,+2 ,+3
Quantum Numbers - VI
Allowed Values
n
1
L
0
0
mL
0
0 -1 0 +1 0 -1 0 +1
2
3
1
0
1
4
2
0
+1/2 -1/2
All + or - 1/2 spin
2
3
0 -1 0 +1
-2 -1 0 +1 +2
ms
1
-2 -1 0 +1 +2
-3 -2 -1 0 +1 +2 +3
Quantum Numbers - VII
Electron Orbitals
Noble Gases
Number of Electrons
1s2
Element
2
He
1s2 2s22p6
10
Ne
1s2 2s22p6 3s23p6
18
Ar
1s2 2s22p6 3s23p6 4s23d104p6
36
Kr
1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6
54
Xe
1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6 6s24f14 5d106p6
86
Rn
1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6 6s24f145d106p6
7s25f146d107p6
118
?
Major Conclusions from Studies of
Orbital Stability - I
1) The Effect of Nuclear Charge (Z) on Orbital Energy
He+ & H have one electron but He+ has 2 protons, thus twice the
attractive force on the electrons:
Ionization Energy for the Two: He+ = - 5250 kJ / mole
H = - 1311 kJ / mole
2) The effect of an additional Electron on Orbital Energy
He has two electrons, whereas He+ has only one, the resultant
repulsion of the electrons in He orbital gives a higher orbital
energy (smaller negative number).
E for He+ = -5250 kJ / mole E for He = -2372 kJ / mole
Major Conclusions from Studies of
Orbital Stability - II
3) The Effect of Inner Electrons on the Energy of an Outer Orbital
The inner electrons (1s) shield the outer electrons (2s) from the full
attractive force of the nucleus, making the 2s orbital higher in energy.
This shielding means that the effective nuclear charge(Zeff), the
nuclear charge an electron actually experiences, is less for an electron
in an outer orbital.
E of H 1s = - 1311 kJ/mol and E of Li 2s = - 520 kJ/mol
4) The Effect of Orbital Shape (L value) on Orbital Energy
Because of their different shapes, a 2s electron is, on the average,
slightly further from the nucleus than the 2p, therefore we would
expect a 2s electron to be attracted less strongly and be higher in
energy. But because the 2s electron also has a small probability of
“penetrating” very close to the nucleus, thus lowering the energy of
the 2s electron, making its energy lower than the 2p electron.
Electron Configuration of
Helium and Lithium
• He
•
•
1s2
n=1
n=1
L=0
L=0
mL = 0
mL = 0
• Li
•
•
•
1s2 2s1
n=1
n=1
n=2
L=0
L=0
L=0
mL = 0 ms = + 1/ 2
mL = 0 ms = - 1/ 2
mL = 0 ms = - 1/ 2
ms = + 1/ 2
ms = - 1/ 2
Orbital Box Diagrams - I
Element Symbol
Hydrogen
Helium
Lithium
Beryllium
Fig. 8.5
Electron
Configuration
H
1s1
He
1s2
Li
1s22s1
Be
H
Be
Orbital Box Diagrams
1s
2s
1s
2s
1s
2s
1s
2s
1s22s2
Hund’s Rule
• For an atom in its ground-state configuration,
all unpaired electrons have the same spin
orientation.
• Therefore electrons tend to occupy all free
orbitals and not pair up, so that their spins all
add up to produce a general vector for the
atom.
Orbital Occupancy for the First 10
Elements, H through Ne
Fig. 8.6
Orbital Box Diagrams - II : B
B (5 e-)
1s2 2s2 2p1
C (6 e-)
1s2 2s2 2p2
N (7 e-)
1s2 2s2 2p3
e-)
1s2 2s2 2p4
O (8
F (9 e-)
Ne (10 e-)
Ne
1s
2s
2px
2py
2pz
1s
2s
2px
2py
2pz
1s
2s
2px
2py
2pz
1s
2s
2px
2py
2pz
1s
2s
2px
2py
2pz
1s
2s
2px
2py
2pz
1s2 2s2 2p5
1s2 2s2 2p6
Valence and Core Electrons
• Valence Electrons - Those electrons outside of a
closed electron shell. These electrons take part in
chemical reactions.
• Core Electrons - The electrons in the closed shells.
They cannot take part in chemical reactions.
• Sodium 11 electrons
•
Valence electrons
[Ne] 3s 1 --- one
•
Core electrons
1s 2 2s 2 2p 6 --- Ten
• Chlorine 17 electrons
•
Valence electrons [Ne] 3s 2 3p 5---- seven
•
Core 2 2s 2 2p 6 ---- Ten
Quantum Numbers and the
Number of Electrons
•
•
•
•
•
•
•
•
•
•
n
L
m
# e-
s
##
==========================================================
1
2
3
4
0
0
1
0
1
2
0
1
0 (1s)
0 (2s)
-1,0,+1 (2p)
0 (3s)
-1,0,+1 (3p)
-2,-1,0,+1,+2(3d)
0 (4s)
-1,0,+1 (3p)
+1/2 - 1/2
+1/2 -1/2
+1/2-1/2
+1/2-1/2
+1/2-1/2
+1/2-1/2
+1/2-1/2
+1/2-1/2
2
2
6
2
6
10
2
6
* Denotes a noble gas !!!
Order of Electron Filling
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
7p
2*
4
10*
12
18*
28
30
36*
Electron Configuration - I
•
•
•
•
•
•
•
•
•
•
1s 1
1s 2
1s2 2s 1
1s2 2s 2
1s2 2s 2 2p 1
1s 2 2s 2 2p 2
1s 2 2s 2 2p 3
1s 2 2s 2 2p 4
1s 2 2s 2 2p 5
1s 2 2s 2 2p 6
H
He
Li
Be
B
C
N
O
F
Ne
[He]
[He] 2s 1
[He] 2s 2
[He] 2s 2 2p 1
[He] 2s 2 2p 2
[He] 2s 2 2p 3
[He] 2s 2 2p 4
[He] 2s 2 2p 5
[He] 2s 2 2p6 = [Ne]
Electron Configuration - II
•
•
•
•
•
•
•
•
Na
Mg
Al
Si
P
S
Cl
Ar
[Ne] 3s 1
[Ne] 3s 2
[Ne] 3s 2 3p 1
[Ne] 3s 2 3p 2
[Ne] 3s 2 3p 3
[Ne] 3s 2 3p 4
[Ne] 3s 2 3p 5
[Ne] 3s 2 3p6 == [Ar]
Condensed Ground-State Electron
Configurations in the First Three
Periods
Fig. 8.7
Orbital Box Diagrams - III
Na
Atomic Number
Element
11 Na
Condensed Electron
Configuration
[He] 3s1
12
13
14
15
16
17
18
Orbital Box
Diagrams(3s&3p)
3s
3px
3py
3pz
3s
3px
3py
3pz
3s
3px
3py
3pz
3s
3px
3py
3pz
3s
3px
3py
3pz
3s
3px
3py
3pz
3s
3px
3py
3pz
Ar
[He] 3s2
Mg
[He] 3s23p1
Al
[He] 3s23p2
Si
[He] 3s23p3
P
[He] 3s23p4
S
Cl
[He] 3s23p5
[He] 3s23p6
Ar
Fig. 8.8
Electron Configuration - III
•
•
•
•
•
•
•
•
•
•
•
•
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
[Ar] 4s 1
[Ar] 4s 2
[Ar] 4s 2 3d 1
[Ar] 4s 2 3d 2
[Ar] 4s 2 3d 3
[Ar] 4s 1 3d 5
[Ar] 4s 2 3d 5
[Ar] 4s 2 3d 6
[Ar] 4s 2 3d 7
[Ar] 4s 2 3d 8
[Ar] 4s 1 3d 10
[Ar] 4s 2 3d 10
Or this order is OK !
[Ar] 3d 1 4s 2
[Ar] 3d 2 4s 2
[Ar] 3d 3 4s 2
Anomalies
to Filling
Either order will be OK !
But it’s normally best to
put the one filling last!!!
Anomalies
to Filling
Orbital Box Diagram - IV : Sc
4s
Z = 21
Z = 22
Zn
3d
4s2 3d1
Sc
Ti
[Ar]
[Ar] 4s 2 3d 2
Z = 23
V
[Ar] 4s 2 3d 3
Z = 24
Cr
[Ar] 4s1 3d 5
Z = 25
Mn
[Ar] 4s 2 3d 5
Z = 26
Fe
[Ar] 4s 2 3d 6
Z = 27
Co
[Ar] 4s 2 3d 7
Z = 28
Ni
[Ar] 4s 2 3d 8
Z = 29
Cu
[Ar] 4s 1 3d 10
Z = 30
Zn
[Ar] 4s 2 3d 10
Electron Configuration - IV
•
•
•
•
•
•
Ga
Ge
As
Se
Br
Kr
[Ar] 4s 2 3d 10 4p 1
[Ar] 4s 2 3d 10 4p 2
[Ar] 4s 2 3d 10 4p 3
[Ar] 4s 2 3d 10 4p 4
[Ar] 4s 2 3d 10 4p 5
[Ar] 4s 2 3d 10 4p 6
=
[Kr]
Electron Configuration - V
•
•
•
•
•
•
•
•
•
•
•
•
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
[Kr] 5s 1
[Kr] 5s 2
[Kr] 5s 24d 1
[Kr] 5s 2 4d 2
[Kr] 5s 1 4d 4
[Kr] 5s 1 4d 5
[Kr] 5s 2 4d 6
[Kr] 5s 1 4d7
[Kr] 5s 1 4d 8
[Kr] 4d 10
[Kr] 5s 1 4d 10
[Kr] 5s 2 4d 10
Anomalies to
Filling
Electron Configuration - VI
•
•
•
•
•
•
In
Sn
Sb
Te
I
Xe
[Kr] 5s 2 4d 10 5p 1
[Kr] 5s 2 4d 10 5p 2
[Kr] 5s 2 4d 10 5p 3
[Kr] 5s 2 4d 10 5p 4
[Kr] 5s 2 4d 10 5p 5
[Kr] 5s 2 4d 10 5p 6
= [Xe]
Electron Configuration - VII
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
Cs
Ba
La
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
[Xe] 6s 1
[Xe] 6s 2
[Xe] 6s2 5d 1
[Xe] 6s 2 5d 1 4f 1
[Xe] 6s 2 4f 3
[Xe] 6s 2 4f 4
[Xe] 6s 2 4f 5
[Xe] 6s 2 4f 6
[Xe] 6s 2 4f 7
[Xe] 6s 2 3d 1 4f 7
[Xe] 6s 2 4f 9
[Xe] 6s 2 4f 10
[Xe] 6s 2 4f 11
[Xe] 6s 2 4f 12
[Xe] 6s 2 4f 13
[Xe] 6s 2 4f 14
[xe] 6s 2 3d 1 4f 14
Anomalies to
Filling
Electron Configuration - VIII
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
[Xe] 6s 2 4f 14 5d 2
[Xe] 6s 2 4f 14 5d 3
[Xe] 6s 2 4f 14 5d 4
Anomalies
[Xe] 6s 2 4f 14 5d 5
Filling
[Xe] 6s 2 4f 14 5d 6
[Xe] 6s 2 4f 14 5d 7
[Xe] 6s 1 4f 14 5d 9
[Xe] 6s 1 4f 14 5d 10
[Xe] 6s 2 4f 14 5d 10
[Xe] 6s 2 4f 14 5d 10 6p 1
[Xe] 6s 2 4f 14 5d 10 6p 2
[Xe] 6s 2 4f 14 5d 10 6p 3
[Xe] 6s 2 4f 14 5d 10 6p 4
[Xe] 6s 2 4f 14 5d 10 6p 5
[xe] 6s 2 4f 14 5d 10 6p 6 = [Rn]
to
Electron Configuration - IX
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
[Rn] 7s 1
[Rn] 7s 2
[Rn] 7s 2 6d 1
[Rn] 7s 2 6d 2
[Rn] 7s 2 5f 2 6d 1
[Rn] 7s 2 5f 3 6d 1
[Rn] 7s 2 5f 4 6d 1
[Rn] 7s 2 5f 6
[Rn] 7s 2 5f 7
[Rn] 7s 2 5f 7 6d 1
[Rn] 7s 2 5f 9
[Rn] 7s 2 5f 10
[Rn] 7s 2 5f 11
[Rn] 7s 2 5f 12
[Rn] 7s 2 5f 13
[Rn] 7s 2 5f 14
[Rn] 7s 2 5f 14 6d 1
Fr
Ra
Ac
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
Anomalies to
Filling
The Periodic Table of the Elements
Electronic Structure
H
Li Be
NaMg
K Ca Sc Ti
Rb Sr Y Zr
Cs Ba La Hf
Fr Ra Ac Rf
B C N
Al Si P
V Cr Mn Fe Co Ni Cu Zn Ga Ge As
NbMo Tc R Rh Pd Ag Cd In Sn Sb
Ta W Re uOs Ir Pt Au Hg Tl Pb Bi
Ha Sg
O
S
Se
Te
Po
He
F Ne
Cl Ar
Br Kr
I Xe
At Rn
Ce Pr Nd PmSm EuGd Tb DyHo Er Tm Yb Lu
Th Pa U Np PuAmCm Bk Cf Es FmMd NoLr
“ S” Orbitals
“ P” Orbitals
“ d” Orbitals
“ f ” Orbitals
The Periodic Table of the Elements
Anomolies to Electron Filling
H
Li Be
NaMg
K Ca Sc Ti
Rb Sr Y Zr
Cs Ba La Hf
Fr Ra Ac Rf
B C N
Al Si P
V Cr Mn Fe Co Ni Cu Zn Ga Ge As
Nd Mo Tc Ru Rh Pd Ag Cd In Sn Sb
Ta W Re Os Ir Pt Au Hg Tl Pb Bi
Du Sg Bo HaMe
O
S
Se
Te
Po
F
Cl
Br
I
At
Ce Pr Nd PmSmEu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np PuAmCm Bk Cf Es FmMd No Lr
Anomalous Electron Filling
He
Ne
Ar
Kr
Xe
Rn
A Periodic Table of
Partial Ground-State
Electron Configurations
Fig. 8.9
Fig. 8.10
Electronic Configuration Ions
• Na 1s 2 2s 2 2p 6 3s 1
Na+
1s 2 2s 2 2p 6
• Mg 1s 2 2s 2 2p 6 3s 2
Mg+2 1s 2 2s 2 2p6
• Al 1s 2 2s 2 2p 6 3s 2 3p 1
Al+3 1s 2 2s 2 2p 6
• O 1s 2 2s 2 2p 4
O- 2 1s 2 2s 2 2p 6
• F 1s 2 2s 2 2p 5
F- 1
1s 2 2s 2 2p 6
• N 1s 2 2s 2 2p 3
N- 3
1s 2 2s 2 2p 6