CH 8: ELECTRON
CONFIGURATION &
PERIODICITY
Vanessa N. Prasad-Permaul
Valencia Community College
CHM 1045
1
Electron Configuration of Atoms
Electron Configuration of an atom: a particular
distribution of electrons among available subshells.
Li 3 electrons: 1s2 2s1
Orbital Diagram: a diagram that shows how the
orbitals of a subshell are occupied by electrons.
Li 3 electrons:
1s 2s
2
Electron Configuration of Atoms
Pauli Exclusion Principle: no two electrons in
an atom can have the same four quantum
Numbers.
Rewritten: An orbital can hold at most two
electrons, and then only if the electrons have
opposite spins.
SUBSHELL
NUMBER OF
ORBITALS
MAXIMUM
NUMBER OF
ELECTRONS
s (l = 0)
1
2
p (l = 1)
3
6
d (l = 2)
5
10
f (l = 3)
7
14
3
Electron Configuration of Atoms
EXAMPLE 8.1
Which one of the following orbital diagrams or electron
configurations are possible and which are impossible,
according to the Pauli Exclusion Principle? Explain:
a.
d. 1s32s1
1s 2s
2p
e. 1s22s12p7
b.
1s 2s
2p
f. 1s22s22p63s23p63d84s2
c.
1s 2s
2p
4
Electron Configuration of Atoms
EXERCISE 8.1
Look at the following orbital diagrams and electron
configurations, which are possible and which are not
according to the Pauli Exclusion Principle? Explain:
a.
d. 1s22s22p4
1s 2s
2p
e. 1s22s42p2
b.
1s 2s
2p
f. 1s22s22p63s23p103d10
c.
1s 2s
2p
5
Electron Configuration of Atoms
The Building-Up Principle
Ground State: The electron configuration
associated with the lowest energy level of the atom.
Na 1s22s22p63s1
Excited State: The electron configuration associated
with an atom the energy levels other than the most
stable (ground state).
Na* 1s22s22p63p1 (emission of a yellow light at 589nm)
Energy
s<p<d<f
6
Electron Configuration of Atoms
Rules of Aufbau Principle:
 Lower n orbitals fill first.
 Each orbital holds two electrons; each with different
ms.
 Half-fill degenerate (same energy level) orbitals
before pairing
electrons. (p, d, & f)
  
3px 3py 3pz
NOT   __
7
Electron Configuration of Atoms
Electron Configuration of Atoms
A mnemonic diagram of the Aufbau Principle
I
n
c
r
e
a
s
i
n
g
E
n
e
r
g
y
8
Electron Configuration of Atoms
Element
Diagram
Configuration
Li (Z = 3)
 
1s 2s
1s2 2s1
Be (Z = 4)
 
1s 2s
1s2 2s2
B (Z = 5)
   __ __
1s 2s 2px 2py 2pz
1s2 2s2 2p1
C (Z = 6)
    __
1s 2s 2px 2py 2pz
1s2 2s2 2p2
9
Electron Configuration of Atoms
Element
Diagram
Configuration
O (Z = 8)
    
1s 2s 2px 2py 2pz
1s2 2s2 2p4
Ne (Z = 10)
    
1s 2s 2px 2py 2pz
1s2 2s2 2p6
S (Z = 16)
        
1s 2s 2px 2py 2pz 3s 3px 3py 3pz
1s2 2s2 2p6 3s2 3p4 or [Ne] 3s2 3p4
abbreviations using the noble gases referred to as a
pseudo-noble gas core.
Valence Electrons: an electron in an atom outside the
noble gas or pseudo-noble-gas core.
10
Electron Configuration of Atoms
Table of Electron Configuration using noble gas core
11
Electron Configuration of Atoms
Table of the Valence-shell configurations of the Elements
12
Electron Configuration of Atoms
The building-up order using the Periodic Table.
13
Electron Configuration of Atoms
EXAMPLE 8.2: Use the Aufbau Principle to obtain the
complete electron configuration for the ground state of
the Gallium atom (Z = 31). Abbreviate with the noble
gas core and what is the valence shell configuration?
Gallium (Ga) Z = 31
Full configuration: 1s22s22p63s23p64s23d104p1
Rearranged by shells: 1s22s22p63s23p63d104s24p1
Abbreviated configuration: [Ar] 3d104s24p1
Valence-shell configuration: 4s24p1
14
Electron Configuration of Atoms
EXERCISE 8.2: Use the Aufbau Principle to obtain the
complete electron configuration for the ground state of
the Manganese atom (Z = 25). Abbreviate with the noble
gas core and what is the valence shell configuration?
15
Electron Configuration of Atoms
EXAMPLE 8.3: What are the configurations for the outer
electrons of :
a. Tellurium Z = 52
[Kr] 5s24d105p4
[Kr] 4d105s25p4
5s25p4
a. Nickel Z = 28
[Ar]4s23d8
[Ar] 3d84s2
3d84s2
16
Electron Configuration of Atoms
EXERCISE 8.3: What are the configurations for the
noble gas and the outer electrons of :
a. Arsenic
b. Bromine
c. Silver
d. Calcium
17
Electron Configuration of Atoms
EXERCISE 8.4: The lead atom has a ground state
configuration of [Xe]4f145d106s26p2. find the period and
group for this element. From it’s position in the periodic
table, classify it as main-group element, a transition
element or an inner transition element.
18
Anomalous Electron Configurations
 19 of the predicted configurations from the
periodic table are wrong
 Largely due to unusual stability of both half-filled and
fully filled subshells
Cr (Z=24)
expected configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d4
     __
4s 3d 3d 3d 3d 3d
actual configuration: 1s2 2s2 2p6 3s2 3p6 4s1 3d5
     
4s 3d 3d 3d 3d 3d
19
Orbital Diagrams of Atoms; Hund’s Rule
Hund’s Rule: the lowest energy arrangement of
electrons in a subshell is obtained by putting
electrons into separate orbitals of the subshell
with the same spin BEFORE pairing the electrons.
*
1s
2s
2p
1s 2s
2p
1s 2s
2p
20
Orbital Diagrams of Atoms; Hund’s Rule
EXAMPLE 8.4: Write the orbital diagram for the ground
state of the iron atom. Z = 26
Electron configuration: 1s22s22p63s23p63d64s2
Noble gas: [Ar] 3d64s2
Valence electron: 3d64s2
Orbital Diagram:
1s 2s
2p
3s
3p
4s
3d
21
Orbital Diagrams of Atoms; Hund’s Rule
EXERCISE 8.5: Write the orbital diagram for the ground
state of the phosphorus atom. Z = 15
Electron configuration:
Noble gas:
Valence electron:
Orbital Diagram:
22
Magnetic Properties of Atoms
Paramagnetic Substance: a substance that is weakly
attracted by a magnetic field this attraction is generally the
result of unpaired electrons
Diamagnetic Substance: a substance tht is not attracted
by a magnetic field or is very slightly repelled by such a
field. This property generally means that the substance has
only paired electrons
23
Periodic Properties
The Periodic Law: When the elements are arranged by
atomic number, their physical and chemical properties
vary periodically.
•Atomic Radius
•Ionization Energy
•Electron Affinity
(important in discussions of chemical bonding)
24
Periodic Properties
Representation of Atomic Radii of the
Main-Group Elements
25
Periodic Properties
Two Factors that primarily determine the size of the
outermost orbital:
•Principle quantum number (n) of the orbital; the larger
the n of the orbital, the larger the size of the orbital.
•The effective nuclear charge acting on an electron in the
orbital; as the effective nuclear charge increases, the size
of the orbital decreases by pulling the electrons inward.
•Effective nuclear charge: the positive charge that an
electron experiences from the nucleus, equal to the
nuclear charge but reduced by any shielding or screening
from any intervening electron distribution.
26
Periodic Properties
EXAMPLE 8.5: Refer to the periodic table use the trends
noted for size of atomic radii to arrange the following in
order of increasing atomic radius: Al, C, Si
C is above Si in Group IVA the radius of C is smaller than
that of Si.
Al and Si are in the same period, going to the right of the
table the radius of Si is smaller than that of Al
C, Si, Al
In order of increasing radius
27
Periodic Properties
EXERCISE 8.6: Using the periodic table, arrange the
following in order of increasing atomic radius: Na, Be, Mg.
28
Periodic Properties
Ionization Energy: the minimum energy needed to
remove the highest-energy (the outermost) electron from
the neutral atom in the gaseous state.
Li (1s22s1)
Li+ (1s2) + e-
•Within a period, values tend to increase with atomic
number  the lowest values are found in Group 1A.
•Elements with the lower ionization energy lose
electrons easily
•Noble gases have high ionization energy
•Generally, as atomic numbers increase, ionization
energy increases
29
Periodic Properties
Trends of First Ionization Energy, Ei
Increase
Increase
30
Higher Ionization Energy, Ei1234…
 Easy to remove an electron from a partially filled valence
shell
 Difficult to remove an electron from a filled valence shell
 Large amount of stability associated with filled s & p
subshells
 Na:
1s2 2s2 2p6 3s1
 Mg:
1s2 2s2 2p6 3s2
 Cl:
1s2 2s2 2p6 3s2 3p5
31
Periodic Properties
Ionization Energy, Ei
 Some exceptions/irregularities to general trend
 Ei Be > Ei B we would expect opposite
 Be 4 e 1s2 2s2
 B 5 e 1s2 2s2 2p1
 2s is closer to nucleus than 2p, Zeff for Be is
stronger
 2s is held more tightly and is harder to
remove
32
Periodic Properties
Ionization Energy, Ei
 Ei N > Ei O we would expect opposite
 N 7e 1s2 2s2 2p3
__ __ __
 O 8e 1s2 2s2 2p4
__ __ __
 Only difference is that an electron is being removed from
a half-filled orbital (N) and one from a filled orbital (O)
 Electrons repel each other and tend to stay as far apart
as possible, electrons that are forced together in a
filled orbital are slightly higher in energy so it is easier
to remove one  O < N
33
Periodic Properties
EXAMPLE 8.6: Using the periodic table, arrange the
following in order of increasing ionization energy: Ar, Se, S.
•Se is below S I Group VIA ionization energy of Se
should be lower than S
•S and Ar are in the same period with Z increasing from
S to Ar the ionization energy of S should be lower
than that of Ar.
• Se > S> Ar
34
Periodic Properties
EXERCISE 8.7: The first ionization energy of the chlorine
atom is 1251 kJ/mol. State which of the following values
would be the more likely ionization energy for the iodine
atom. Explain.
a. 1000kJ/mol
or
b. 1400kJ/mol
35
Ionic Radii or size
 Atoms expand when converted to anions
 III A
ns2 np1
__ __ __
 IV A
ns2 np2
__ __ __
 VA
 VI A
 VII A
ns2 np3
ns2 np4
ns2 np5
__ __ __
__ __ __
__ __ __
Adding one electron to each of these will not
add another shell it will just fill an already
occupied p subshell
 Therefore the expansion is due to the
decrease in Zeff and the increase in the
electron-electron repulsions
36
Ionic Radii or size
 Atoms contract when an electron is removed
to form a cation.
 Dec. # of shells
 Inc. Zeff : Less electrons, less shielding, outer
electrons more attracted to nucleus, therefore
smaller more compact
37
Higher Ionization Energy, Ei1234…
 Ionization is not limited to one electron
M + Energy  M+ + e Ei1
M+ + Energy  M2+ + e Ei2
M2+ + Energy  M3+ + eEi3
 Larger amounts of energy are needed for each
successive ionization, harder to remove an
electron from a positively charged cation
38
Periodic Properties
Electron Affinity, Eea
 Energy change that occurs when an electron is
added to an isolated atom in the gaseous state.
 The more negative the Eea , the greater the
tendency of the atom to accept an electron
 Group 7A (halogens) have the most negative Eea,
high Zeff and room in valence shell
 Group 2A and 8A have near zero or slightly
positive Eea
39
Periodic Properties
EXERCISE 8.8: Using the general comments that were
discussed in this section, decide which has the larger
negative electron affinity: C or F.
40
Periodicity in the Main-Group Elements
Alkali Metals
 Group 1A (ns1)





Metallic
Soft
Good Conductors
Low melting point
Lose 1 electron in redox reactions; powerful
reducing agent
 Very reactive
 Not found in elemental state in nature
41
Periodicity in the Main-Group Elements
Alkaline Earth Metals
 Group 2A (ns2)
 Harder, but still relatively soft
 Silvery
 High melting point than group 1A
 Less reactive than group 1A
 Loses 2e- in redox reaction; powerful reducing
agent
 Not found in elemental form in nature
42
Periodicity in the Main-Group Elements
Group 3A (ns2np1)
 All but Boron which is a metalloid
 Silvery
 Good conductor
 Relatively soft
 Less reactive than 1A & 2A
 metals
43
Periodicity in the Main-Group Elements
Halogens
 Group 7A (ns2np5)
 Non-metals
 Diatomic molecules
 Tend to gain e- during redox reaction.
44
Periodicity in the Main-Group Elements
Noble Gases
 Group 8A (ns2np6)
 Colorless, odorless, unreactive gases
 Stable because of the filled subshell
 Makes it difficult to add electrons or remove
electrons
45
Example 1: Electron Config. And NG Abb.
1. Sodium
2. Titanium
3. Argon
46
Example 2: Ionic Radii
Which of the following in each pair has a larger atomic
radius?
1. Carbon or Fluorine
2. Chlorine or Iodine
3. Sodium or Magnesium
4. O or O25. Ca or Ca2+
47
Example 3: Quantum Numbers and Electron Configuration
What are the 4 quantum numbers for the following?
Remember you are only interested in the last
electron!!
1.
C
2.
Na+
3.
S
4.
N348
Example 4: Electron config. and NG Abb.
1. Cl-
2. F-
3. Ca2+
4. Na+
49