George Mason University
General Chemistry 211
Chapter 8
Electron Configuration and Chemical Periodicity
Acknowledgements
Course Text: Chemistry: the Molecular Nature of Matter and
Change, 7th edition, 2011, McGraw-Hill
Martin S. Silberberg & Patricia Amateis
1/13/2015
The Chemistry 211/212 General Chemistry courses taught at
George Mason are intended for those students enrolled in a science
/engineering oriented curricula, with particular emphasis on
chemistry, biochemistry, and biology The material on these slides is
taken primarily from the course text but the instructor has modified,
condensed, or otherwise reorganized selected material.
Additional material from other sources may also be included.
Interpretation of course material to clarify concepts and solutions to
problems is the sole responsibility of this instructor.
1
Electron Configuration & Chemical Periodicity



Development of the Periodic Table
Characteristics of Many-Electron Atoms
 The Electron Spin Quantum Number
 The Pauli Exclusion Principle
 Electrostatic Effects and Energy-Level Splitting
Development of the Quantum Mechanical Model of the
Periodic Table
 Building up of Periods 1 & 2
 Building up of Period 3
 Electron Configuration Within Groups
 Building up Period 4
 General Principles of Electron Configuration
 Unusual Configurations: Transition and Inner Transition
Elements
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2
Electron Configuration & Chemical Periodicity

Trends in Three Key Atomic Properties
 Trends in Atomic Size
 Trends in Ionization Energy (IA)
 Trends in Electron Affinity (EA)

Atomic Structure and Chemical Reactivity
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3
Electron Configuration

The work of Balmer (Rydberg equation) and Bohr (Bohr
Postulates) invoked the idea that spectra lines of
compounds represented different wavelengths (Balmer)
and energy levels (Bohr) of the radiation.
 1 1 
1
= R 2 - 2
λ
 2 n2 
 1 1 
E = hν = - 2.18  10-18 J  2 - 2 
n n 
i 
 f

Each energy level was designated by the whole number
integer, n.

“n” could have any value from 1 -

The higher the value of “n,” the smaller the wavelength,
which equates to higher frequency and higher energy
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 (infinity)
4
Electron Configuration

The work of Balmer, Bohr, and Einstein,
de Broglie, Heisenberg and many others lead to the
development of Quantum Mechanics

Quantum mechanics, also known as quantum physics or
quantum theory, is a branch of physics providing a
mathematical description of the wave-particle duality of
matter and energy

In quantum chemistry each atom is distinguished by its
unique number of electrons, which is matched by an equal
number of protons in the nucleus (atomic number, Z)

Each new element has one more electron than its
predecessor
Hydrogen (H) 1 e-; Helium (He) 2 e-
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5
Electronic Configuration

The electrons are configured (distributed) into
“orbitals,” which represent energy levels

Electrons are able to move from one orbital (energy
level) to another by emission or absorption of a
quantum of energy, in the form of a photon

Knowledge of the electron configuration of different
atoms is useful in understanding the structure of the
periodic table of elements

The concept is also useful for describing the chemical
bonds that hold atoms together

In bulk materials this same idea helps explain the
peculiar properties of lasers and semiconductors
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6
Electron Configuration

According to quantum mechanics each electron is
described by 4 Quantum numbers
 Principal Quantum Number
(n)
 Angular Momentum Quantum Number
(l)
 Magnetic Quantum Number
(ml)
 Spin Quantum Number
(ms)
The first 3 quantum numbers define the wave
function of the electron’s atomic orbital, i.e., it size
and general energy level
The fourth quantum number refers to the
Spin Orientation
of the 2 electrons that occupy an atomic orbital
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7
Electronic Configuration

Quantum Numbers and Atomic Orbitals
 The Principal Quantum Number (n) represents the
“Shell Number” in which an electron “resides”
● It represents the relative size of the orbital
● Equivalent to periodic chart Period Number
● Defines the principal energy of the electron
● The smaller “n” is, the smaller the orbital size
● The smaller “n” is, the lower the electron energy
● n can have any positive value from
1, 2, 3, 4 … 
(Currently, n = 7 is the maximum known)
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8
Electronic Configuration

Quantum Numbers and Atomic Orbitals (Con’t)
 The Angular Momentum Quantum Number (l)
distinguishes “sub shells” within a given shell
● Each main “shell,” designated by quantum
number “n,” is subdivided into:
l = n - 1 “sub shells”
● (l) can have any integer value from 0 to n - 1
● The different “l” values correspond to the
s, p, d, f designations used in the electronic
configuration of the elements
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Letter
s
p
d
f
(g)
l value
0
1
2
3
(4)
9
Electronic Configuration

Quantum Numbers and Atomic Orbitals (Con’t)
 The Magnetic Quantum Number (ml) defines the
atomic orbitals within a given sub-shell
● Each value of the angular momentum number (l)
determines the number of atomic orbitals
● For a given value of “l,” ml can have any integer
value from –l to +l
ml = –l to +l
● Each orbital has a different shape and orientation
(x, y, z) in space
● Each orbital within a given angular momentum
number sub shell (l) has the same energy
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10
Electron Configuration

Quantum Numbers and Atomic Orbitals (Con’t)
 The Spin Quantum Number (ms) refers to the two
possible spin orientations of the electrons residing
within a given atomic orbital
● Each atomic orbital can hold only:
two (2) electrons
● Each electron has a “spin” orientation value
● The spin values must oppose one another
● The possible values of ms spin values are:
+1/2 and –1/2
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11
Stern-Gerlach Experiment
A beam of H atoms can be separated into 2 beams of
opposite electron spin in a magnetic field
ms ( –1/2 ) electrons have a slightly greater energy than
ms ( +1/2 ) electrons
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12
Representation of electron spin
A spinning charged particle aligns in a magnetic field
depending on spin state
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13
Summary of Quantum Numbers
Name
Symbol
Permitted
Values
positive integers
(1, 2, 3, …)
principal
n
angular
momentum
l
integers from
0  n -1
magnetic
ml
integers from
–l  0  +l
spin
ms
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Property
orbital energy
(size)
orbital shape
The l values
0, 1, 2, and 3
correspond to
s, p, d, f
orbitals, respectively
orbital orientation
+1/2  or – 1/2  direction of e- spin
14
Summary of Quantum Numbers
Name,
Symbol
(Property)
Allowed Values
Principal, n
(size, energy)
Positive integer
(1, 2, 3, ...)
Angular
momentum, l
(shape)
Magnetic, ml
(orientation)
Quantum Numbers
n=2
n=1
n=3
l = 0  n -1
0(s), 1(p),
2(d), 3(f)
l=0
(1s)
l=0
(2s)
l=1
(2p)
-l,…,0,…,+l
0
0
-1 0 +1
l=0
(3s)
l=2
(3d)
0 -1 0 +1
-2
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l=1
(3p)
-1
0 +1 +2
15
Electron Configuration

An electron configuration of an atom is a particular
distribution of electrons among available sub shells
 The configuration notation lists the subshell symbols (s,
p, d, f…) sequentially with a superscript indicating the
number of electrons occupying that subshell
 Ex: lithium (Period (n) = 2, Atomic No 3) has
2 electrons
in the “1s” sub shell
1 electron
in the “2s” sub shell
1s2 2s1
Fluorine (Period (n) 2, Atomic No 9) has
1/13/2015
2 electrons in the
“1s” sub shell
2 electrons in the
“2s” sub shell
5 electrons in the
“2p” subshell
1s2 2s2 2p5
16
Electron Configuration



A unique set of the first 3 quantum numbers (n, l, m l)
defines an “Orbital”
An orbital can contain a maximum of 2 electrons, each
with a different “spin” (+1/2 or -1/2)
An orbital diagram is notation used to show how the
orbitals of a sub shell are occupied by electrons
 Each orbital is represented by a circle
 Each orbital can have a maximum of 2 electrons
 Each group of orbitals is labeled by its
Sub Shell Notation (s, p, d, f)
 Electrons are represented by arrows:
up () for ms = +1/2 and down () for ms = -1/2
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1s
2s
2p
17
The Pauli Exclusion Principle

The Pauli Exclusion Principle
 No two electrons in an atom can have the
same four quantum numbers
 An orbital (unique combination of n, l, m l )
can hold, at most, two electrons
 Two electrons in the same Orbital have
opposite spins
+1/2 
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-1/2 
18
The Pauli Exclusion Principle

The maximum number of electrons and their orbital
diagrams are:
Sub
Shell
No.
Orbitals
Values
(-l to +l)
Max No.
Electrons
s (l = 0)
1
(0)
2
p (l = 1)
3
(-1,0,+1)
6
d (l =2)
5
(-2,-1,0,+1,+2)
10
f (l =3)
7 (-3,-2,-1,0,+1,+2,+3)
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14
19
The Pauli Exclusion Principle
n l (n-1)
1 0 (1s)
ml
0
2 0 (2s)
1 (2p)
0
-1 0 +1
3 0 (3s)
1 (3p)
2 (3d)
0
-1 0 +1
-2 -1 0 +1 +2
4 0 (4s)
1 (4p)
2 (4d)
3 (4f)
0
1 0 +1
-2 -1 0 +1 +2
-3 -2 -1 0 +1+2 +3
5 0 (5s)
0
1 (5p)
-1 0 +1
2 (5d)
-2 -1 0 +1 +2
3 (5f)
-3 -2 -1 0 +1+2 +3
4 (5g) -4 -3 -2 -1 0 +1 +2 +3 +4
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20
Noble Gas Electronic Configurations

In the following slides electronic configurations of the
elements in the Periodic Table will be discussed

Electronic configurations can become quite complex
(lengthy) as the Atomic Number increases

A condensed form of the Electronic Configuration of a
given element or ion is often used

A symbol, [X], representing the electron configuration of
the Noble gas in the period just above the element of
interest is substituted for the detail configuration

The following slide illustrates the Noble Gas configurations
and the “Condensed Form” symbol used with other
elements
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21
Configurations and the
Periodic Table
Condensed Electronic Configurations
Full Electronic Configuration
Condensed
Electronic
Configuration
Helium
1s2
[He] tot e- 2
Neon
1s22s22p6
[Ne] tot e- 10
Argon
1s22s22p63s23p6
[Ar]
tot e- 18
Krypton
1s22s22p63s23p63d104s24p6
[Kr]
tot e- 36
Xenon
1s22s22p63s23p63d104s24p64d105s25p6
[Ze]
tot e- 54
Beryllium
1s22s2
[He] 2s2 tot e- 4
Magnesium 1s22s22p63s2
[Ne] 3s2 tot e- 12
Calcium
1s22s22p63s23p64s2
[Ar] 4s2 tot e- 20
Sodium Ion
(Na) 1s22s22p63s1  (Na+) 1s22s22p6 + 1e-
[Ne] + 1e-
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22
Electron Configuration

Quantum Number
n = 1 (Period 1)
l values = 0 to (n-1) = 0 to (1 -1) = 0  l = 0 (s
orbital)
ml values = -l,…0,…+l
= 0 (1 s orbital)
ms values = -1/2 & +1/2 = (2 e- per orbital)
1s orbital
Z = 1 Hydrogen 1s1
Z = 2 Helium
1s2
Thus, for n = 1 there is one orbital (s) which can
accommodate 2 elements – Hydrogen & Helium
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23
Electron Configuration

Quantum Number n = 2 (Period 2)
l values = 0 to (n-1) = 0 to (2-1) = 0 to 1  l = 0(s), 1(p)
For l = 0 (s) ml = -l 0 + l = 0 (one 2s orbital, 2 electrons)
ms values = -1/2 & +1/2
For l = 1 (p) ml = -1 0 +1 (three 2p orbitals, 6 electrons)
ms values = -1/2 & +1/2 in each orbital
2s orbitals
Z=3
Lithium
1s22s1
or [He]2s1
Z=4
Beryllium
1s22s2
or [He]2s2
Z=5
Boron
1s22s22p1 or [He]2s22p1
Z=6
Carbon
1s22s22p2 or [He]2s22p2
Z=7
Nitrogen
1s22s22p3 or [He]2s22p3
Z=8
Oxygen
1s22s22p4 or [He]2s22p4
Z=9
Fluorine
1s22s22p5 or [He]2s22p5
Z=10
Neon
1s22s22p6 or [He]2s22p6
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2p orbitals
24
Electron Configuration

With Sodium (Z = 11), the 3s sub shell begins to fill
Z=11 Sodium
1s22s22p63s1 or [Ne]3s1
Z=12 Magnesium 1s22s22p63s2 or [Ne]3s2
Starting with Z = 13, the 3p sub shell begins to fill
Z=13 Aluminum 1s22s22p63s23p1 or [Ne]3s23p1
Z=18
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Argon 1s22s22p63s23p6
or [Ne]3s23p6
25
Electron Configuration

Electrostatic Effects and Energy-Level Splitting
 The principal quantum number (n) defines the
energy level of an atom
● The higher the “n” value, the higher the energy
level
 The unique values of the principal quantum
numbers of multi-electron atoms (n, l, ml) define a
unique energy level for the orbital of a given
electron
● The energy of a given orbital depends mostly on
the value of the principal quantum number (n),
i.e. its size, and to a lesser degree on the shape
of the orbital represented by the various values
of the magnetic quantum number (l)
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26
Electron Configuration

The energy states of multi-electron atoms arise from
2 counteracting forces:
 Nucleus – Positive protons attract Negative electrons
 Electron – Negative electrons repulse each other
 Nuclear protons create a pull (attraction) on electrons
 Higher nuclear charge (Z) lowers orbital energy
(stabilizes system) by increasing proton-electron
attractions
● The energy required to remove the 1s electron from
Hydrogen (H), Z =1, is much less than the energy to
remove the 1s electron from the Li2+ ion, Z = 3
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27
Electron Configuration

Effect of Nuclear Charge (Z) on Orbital Energy
Greater Nuclear Charge lowers orbital energy
making it more difficult to remove the electron
from orbit
The absolute value of the 1s orbital energy is related
directly to Z2
Energy required to remove 1s electron from H
1311 kJ/mol (Z= +1, Least stable)
Energy required to remove 1s electron from He+
5250 kJ/mol (Z = +2)
Energy required to remove 1s electron from Li+
11815 kJ/mol (Z = +3, Most stable)
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28
Electron Configuration
 Shielding – Effect of Electron Repulsions on Orbital Energy
 Electrons feel repulsion from other electrons somewhat
shielding (counteracting) the attraction of the nuclear
protons
 Shielding (screening) lowers the full nuclear charge to
an “Effective Nuclear Charge (Zeff)
 The lower the Effective Nuclear Charge, the easier it is
to remove an electron
● It takes less than half as much energy to remove an
electron from Helium (He) (2373 kJ/mol) than from
He+ (5250 kJ/mol) because the second electron in He
repels the first electron and effectively shields the
first electron from the full nuclear charge (lower Zeff)
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29
Electron Configuration

Penetration: Effects of orbital shape
 The shape of an atomic orbital affects how close an
electron moves closer to nucleus, i.e., the level of
penetration
 Penetration and the resulting effects of shielding on a
atomic orbital causes the energy level (n) to be split
into sublevels of differing energy representing the
various values of the magnetic quantum number (l)
 The lower the value of the magnetic quantum number
(l), the more its electrons penetrate
Order of Sublevel Energies
s (l=0) < p(l=1) < d(l=2) < f(l=3)
 Each of the orbitals for a given value of l
(ml = -l 0 +l) has the same energy
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30
Aufbau Principle

Aufbau Principle – scheme used to reproduce
the ground state electron configurations of atoms
by following the “building up” order based on
relative energy levels of quantum subshells

The “building up” order corresponds for the most
part to increasing energy of the subshells

By filling orbitals of the lowest energy first, you
usually get the lowest total energy (“ground
state”) of the atom
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31
Aufbau Principle

Listed below is the order in which all the possible subshells fill with electrons
Note the order does NOT follow the strict numerical
subshell order shown on slide 20
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f
You need not memorize this order
The next slide provides a pictorial providing an easier way
of the viewing the ‘build-up” order
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32
Order for Filling Atomic Subshells
Principal
Quantum
No. (n)
n=1
n=2
n=3
n=4
n=5
n=6
n=7
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Angular Momentum (l)
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
• Setup rows for each
Principal Quantum No. (n)
• Set columns for each
Angular Momentum (l)
3d
4d
5d
6d
7d
• Draw a series of diagonals
4f
5f
6f
7f
• Order of filling is the order
in which diagonals strike
subshells
• Note the 4s subshell is
filled before the 3d
subshell because the 4s
electrons
are at lower energy levels
than the 3d electrons
33
Aufbau Principle

Every atom has an infinite number of possible electron
configurations (electrons can be raised to any number of
energy (n) levels)
 The configuration associated with the lowest energy
level of the atom is called the
“ground state”
 Other configurations correspond to
“excited states”
 Tables on the next 3 slides list the ground- state
configurations of atoms up to krypton
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34
Partial Orbital Diagrams
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35
Partial Orbital Diagrams

Chromium (Cr) relative to Vanadium (V)

The Cr 4s1 subshell is filled before the 3d subshell is completed

An [Ar]3d44s2 orbital configuration would be expected for
ground state Cr, but the [Ar]3d54s1 orbital is lower in energy
Partial Orbital Diagrams and Electron Configurations* for the Elements in Period 4
Cr
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36
Partial Orbital Diagrams

Copper (Cu) relative to Nickel (Ni)

Copper would be expected to have a ground state
configuration of [Ar]3d94s2

The [Ar]3d104s1 configuration is actually lower in energy
Partial Orbital Diagrams and Electron Configurations* for the Elements in Period 4
Ni
Cu
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37
Orbital Energy Levels in
Multi-Electron Systems
3d
4s
Energy
3p
3s
2p
2s
1s
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3d orbitals would be
expected to be filled before
4s orbitals
Actual order of filling
depends on total ground
state energy of the atom
3d and 4s orbitals are very
close in energy
Selected 4s, 5s, 6s, 7s levels
are filled before 3d, 4d, 4f,
5f, respectively
(see slides 31 & 32)
38
Configurations and the
Periodic Table

Electrons that reside in the outermost shell of an atom
- or in other words, those electrons outside the “noble gas
core” - are called valence electrons
 These electrons are primarily involved in chemical
reactions
 Elements within a given group have the same valence
shell configuration
 This accounts for the similarity of the chemical
properties among groups of elements
n
n
n
n
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=
=
=
=
2
3
4
5
Li
Na
K
Rb
–
–
–
–
2s1
3s1
4s1
5s1
Be
Mg
Ca
Sr
–
–
–
–
2s2
3s2
4s2
5s2
39
Configurations and the
Periodic Table
 Noble gas core: an inner shell configuration resembling
one of the noble gases (He, Ne, Ar, Kr, Xn)
 Pseudo-noble gas core: noble gas core + (n-1)d10
electrons: Ex Sn  Sn+4
Sn ([Kr] 5s2 4d10 5p2)  Sn+4 ([Kr] 4d10 + 4 e-
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40
Configurations and the
Periodic Table

Configurations of Main Group Ions
 Noble gases have filled outer energy levels (ns2np6),
have very high Ionization Energies (IEs), and positive
(endothermic) Electron Affinities (EAs); thus do not
readily form ions
 Elements in Groups 1A, 2A, 6A, 7A that readily form
ions by gaining electrons (1A & 2A) or losing electrons
(6A & 7A) attain a filled outer level conforming to a
Noble Gas configuration
 Such ions are said to be “Isoelectronic” with the nearest
Noble gas configuration
Na (1s22s22p63s1)  Na+ (1s22s22p6) + 1eIsoelectronic with [Ne] + 1e-
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41
Configurations and the
Periodic Table

The energy needed to remove the electrons from metals in
groups 1A, 2A, 6A, 7A, is supplied during exothermic
reactions with nonmetals

Attempts to remove more than 1 electron from group 1A
or 2 electrons from group 2A metals would mean removing
core (not valence) electrons requiring significantly more
energy than is available from a reaction with a non-metal
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42
Configurations and the
Periodic Table

The larger metals from Groups 3A, 4A, and 5A form
cations through a different process

It would be energetically impossible for them to lose
enough electrons to attain a noble gas configuration
Ex: Tin (Sn), Z = 50 would have to lose 14 electrons (two
5p, ten 4d, two 5s) to be isoelectronic with Krypton: Z =36

Instead, tin loses fewer electrons and still attains one or
more stable pseudo-noble gas configurations
Sn ([Kr] 5s24d105p2)  Sn4+ ([Kr] 4d10) + 4e-
Stability comes from empty 5s & 5p sublevels
and a filled inner 4d sublevel (n-1)d10 configuration
Pseudo-Noble Gas Configuration
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43
Practice Problem
Which of the following electron configurations represents an
excited state?
a. He: 1s2
b. Ne: 1s2 2s2 2p6
c. Na: 1s2 2s2 2p6 3s1
d. P: 1s2 2s2 2p6 3s2 3p2 4s1
e. N: 1s2 2s2 2p3
Ans: d
Ground state for Phosphorus is:
1s2 2s2 2p6 3s2 3p3
The 3p subshell would continue to fill before
the 4s subshell would start to fill
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44
Practice Problem
What is the electron configuration for the valence electrons
of Technetium (Tc, Z = 43)?
a. 4d55s2
b. 5s25d4
d. 4d65s2
Ans: a
e. 3d44s2
4d55s2
c. 4s24d4
5 + 2 = 7 valence electrons
Technetium (atomic no. = 43 = 43 total electrons)
Select “Noble Gas” Configuration prior to Technetium (Kr)
1s22s22p63s23p64s23d104p6  [Kr] 36 e[Kr] + 4d55s2 = 36 + 7 = 43 = Technetium
Note: 4d orbitals filled before 5p orbitals (Aufbau)
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45
Practice Problem
What is the electron configuration for the valence electrons
of Polonium (Po, Z=84)?
a. 6s26p2
b. 6s25d106p4
c. 6s25d106p6
d. 6s26p4
e. 7s26p4
Ans: b
6s25d106p4
2 + 10 + 4 = 16 valence electrons
Polonium (atomic no. 84 = 84 total electrons)
Select “Noble Gas” Configuration prior to Polonium Ze(54)
Xenon 1s22s22p63s23p63d104s24p64d105s25p6 [Ze] 54 e84 – 54 = 30 electrons which must include 14 electrons
that fill in the 4f orbitals that start with Lanthanum
30 -14 = 16 – 10 (filled 5d10) = 6 = 6s25d106p4
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46
Configurations and the
Periodic Table

The following slide illustrates how the periodic table
provides a sound way to remember the Aufbau sequence
 In many cases you need only the configuration of the
outer electrons
 You can determine this from their position on the
periodic table
 The total number of valence electrons for an atom
equals its group (vertical column) number
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47
Periodic Table (Subshells)
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48
Configurations and the
Periodic Table
s block
Main block = s + p blocks
p block
d block
Transition Elements
f block

Inner Transition
Elements 
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49
Orbital Diagrams

Consider carbon (Z = 6) with the ground state
configuration 1s22s22p2
 Three possible arrangements are given in the following
orbital diagrams.
1s
2s
2p
Diagram 1:
Diagram 2:
Diagram 3:

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Each state has a different energy and
different magnetic characteristics
50
Orbital Diagrams

Hund’s rule states that the lowest energy arrangement
(the “ground state”) of electrons in a sub-shell is obtained
by putting electrons into separate orbitals of the sub shell
with the same spin before pairing electrons
 Looking at carbon again, we see that the ground state
configuration corresponds to diagram 1 when following
Hund’s rule
1s
2s
2p
 Note: The 2 e- in the 2p orbitals are shown as “up”
arrows representing the +1/2 spin state, which has
lower energy the -1/2 spin state
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51
Orbital Diagrams

To apply Hund’s rule to Oxygen, whose ground state
configuration is 1s22s22p4, place the first seven electrons
as follows
1s

2p
The last electron is paired with one of the 2p
electrons to give a doubly occupied orbital, i.e.,
a +½ spin state and a – ½ spin state
1s
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2s
2s
2p
52
Summary

Pauli Exclusion principle: no 2 e-s in an atom can have the
same four quantum numbers

Aufbau Principle: obtain electron configurations of the
ground state of atoms by successively filling subshells with
electrons in a specific order

Hunds 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 paring
them
Recall: +1/2 spin has lower energy then -1/2 spin
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53
Periodic Properties

Two factors determine the size of an atom
 One factor is the principal quantum number, n.
The larger “n” is , the larger the size of the orbital
 The other factor is the effective nuclear charge
(slide 28), which is the positive charge an electron
experiences from the nucleus minus any “shielding
effects” from intervening electrons

The Periodic Law states that:
When the elements are arranged by atomic
number, their physical and chemical properties
vary periodically – across the periodic chart row
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54
Periodic Properties – Atomic Size

Atomic Size, Ionization Energy, Electron Affinity
 Atomic radius
● Within each Period (across horizontal row), the
atomic radius tends to decrease with increasing
atomic number (nuclear charge more dominant
than electron repulsion)
● Within each Group (down a vertical column), the
atomic radius tends to increase with increasing
period number (electron repulsion dominates
nuclear charge increase)
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55
Periodic Properties – Atomic Size

Representation of atomic radii (covalent radii) of the maingroup elements (neutral atoms)
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56
Periodic Properties – Atomic Size
Elements vs Ions


Ionic Size increase
down a group
 Number of energy
levels increases
Ionic Size becomes
more complicated
across a period
 Decreases among
cations
 Increase
dramatically with
first anion
 Decreases within
anions
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57
Periodic Properties – Atomic Size

Ionic Size and Atomic Size
 Cations are smaller than their parent atoms
● Electrons are removed from the outer level
● Resulting decrease in electron repulsions
allows nuclear charge to pull remaining
electrons closer
 Anions are larger than their parent atoms
● Electrons added to outer level
● Resulting in increased electron repulsion
allowing them to occupy more space
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58
Periodic Properties – Ionization Energy
 Ionization energy
● The first ionization energy of an atom is
the minimal energy needed to remove the
highest energy (outermost) electron from the
neutral atom
● For a Lithium atom, the first ionization energy
is illustrated by:
Li(1s22s1) → Li+(1s2) + e-
IE = 520 kJ/mol
Endothermic (requires energy input)
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59
Periodic Properties – Ionization Energy
 Ionization energy (IE)
 There is a general trend that ionization energies
increase with atomic number within a given
period
 This follows the trend in size, as it is more
difficult to remove an electron that is closer to
the nucleus
 For the same reason, we find that ionization
energies, again following the trend in size,
decrease descending down a column of elements
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60
Ionization Energy vs Atomic Number
Noble gases have highest IE’s
Alkali metals have lowest IE’s
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61
Periodic Properties – Ionization Energy
Successive Ionization Energies of the First Ten Elements (kJ / mol*
Ionization Energies to the “Right” of the a vertical
line correspond to removal of electrons from the
“Core” of the atom
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62
Periodic Properties – Ionization Energy
 Ionization energy (IE)
● The electrons of an atom can be removed
successively
 The energies required at each step are known as
the first ionization energy, the second
ionization energy, and so forth
 Successive Ionization Energies increase because
each electron is pulled away from an ion with a
progressively higher positive charge, i.e., a more
effective nuclear charge
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Exceptions to Ionization Energy Trends

A IIIA element , such as Boron (2s22p1), has a smaller
ionization energy (IE) than the preceding IIA element
Beryllium (2s2) because one np electron is more easily
removed than the second ns electron

A VIA element, such as oxygen (2s22p4), has smaller
ionization Energy than the preceding VA element nitrogen
(2s22p3). As a result of repulsion it is easier to remove an
electron from the doubly occupied 2p orbital of the VI
element that from a singly occupied p orbital of the
preceding VA element
Nitrogen 2s22p3
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Oxygen 2s22p4
64
Periodic Properties – Electron Affinity
 Electron Affinity (EA): the energy change for the
process of adding an electron to a neutral atom in the
gaseous state to form a negative ion, i.e., an Anion
● 1st Electron Affinity – Formation of 1 mole of
monovalent (1-) gaseous ions
Atoms(g) + e-  ion-(g) E = EA1
● For the formation of the Chloride ion (Cl-) from the
Chlorine atom, the first electron affinity is illustrated
by:
Cl([Ne]3s 3p ) + e  Cl ([Ne]3s 3p )
2
5
-
-
2
6
Electron Affinity = EA1 = 349 kJ/mol
Exothermic (releases energy)
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65
Periodic Properties – Electron Affinity
 Electron Affinity (EA)
● The more negative the electron affinity, the more
stable the negative ion that is formed
● Broadly speaking, the general trend goes from
lower left to upper right as electron affinities
become more negative
● Highest electron affinities occur for halogens, F
and Cl
● Negative values indicate that energy is released
when the Anion forms
● Note: Electron Affinity is not the same as
Electronegativity – relative ability of a bonded
atom to attract shared electrons
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Periodic Properties – Electron Affinity
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67
Periodic Properties – Atomic Size


Atomic Size

Atomic Size (neutral atoms & ions) increases down a main group

Atomic Size (neutral atoms & ions) decreases across a Period

Atomic Size remains relatively constant across a transition series
Ionization Energy
 First Ionization Energy (remove outermost e-) is inversely related to
atomic size
 1st Ionization Energy decreases down a group
 1st Ionization Energy increases across period
 Successive IEs show very large increases after 1st electron is
removed

Electron Affinity
 Similar patterns (with many exceptions) to ionization Energy (lower
left to upper right)
 Highest electron affinities occur for halogens, F and Cl
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Periodic Properties - Summary
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69
Atomic Structure / Chemical Reactivity

Metals
 Metals are located in the left and lower threequarters of the Periodic Table
 Typical Properties
● Shiny Solids
● High Melting Points
● Good Thermal & Electrical Conductors
● Malleable – Drawn into wires and rolled into
sheets
● Lose electrons to non-metals
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70
Atomic Structure / Chemical Reactivity

Non-Metals
 Non-metals are located in the upper right quarter
of the Periodic Table
 Not Shiny
 Low Melting Points
 Poor Thermal & Electrical Conductors
 Crumbly Solids or gases
 Gain Electrons from Metals
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Atomic Structure / Chemical Reactivity

Metalloids (semi-metals)
 Located between Metals & Non-Metals in the
Periodic Chart
boron, silicon, germanium, arsenic,
antimony, tellurium, and polonium
 An element that exhibits the external characteristics
of a metal, but behaves chemically more as a
nonmetal
 Arsenic, for example, is a metalloid that has the
visual appearance of a metal, but is a poor
conductor of electricity
 The intermediate conductivity of metalloids means
they tend to make good semiconductors
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Atomic Structure / Chemical Reactivity

Metalloids (semi-metals)
 The electronegativities and ionization energies of
the metalloids are between those of the metals and
nonmetals, so the metalloids exhibit characteristics
of both classes
 The reactivity of the metalloids depends on the
element with which they are reacting
Ex. Boron
● Acts as a nonmetal when reacting with Sodium
● Acts as a metal when reacting with Fluorine
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Atomic Structure / Chemical Reactivity

Metalloids (semi-metals)
 The boiling points, melting points, and densities of
the metalloids vary widely
 As a rule, metalloids do not form multiple bonds
 Compounds containing these elements will often
show an incomplete octet around the central atom
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74
Atomic Structure / Chemical Reactivity

Metallic Behavior decrease from left to right and increases
from top to bottom in Periodic Tables
Metals
Non-Metals
Transition Elements
f-block Inner Transition Elements
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Atomic Structure / Chemical Reactivity

Metallic Behavior
 Metals tend to “Lose” electrons
 Metals tend to lose electrons during chemical
reactions because they have “Low” ionization
energies compared to non-metals
 Elements generally tend to increase their metallic
character going down a Periodic Table group
 The greatest contrast in changing metallic character
is in groups 3A – 6A
● Elements at the top tend to form “Anions”, i.e.,
more non-metallic character, while those at the
bottom tend to form metallic “Cations”
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Atomic Structure / Chemical Reactivity

Metallic Behavior (Con’t)
 Nitrogen (N) & Phosphorus (P), both non-metals
tend to form 3- anions
 Arsenic (As) (period 4) & Antimony (Sb) (period 5)
are metalloids and generally do not form ions
 Bismuth (Bi) (period 6) is a typical metal forming
mostly ionic compounds as a 3+ cation
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77
Atomic Structure / Chemical Reactivity

Metal Behavior (Con’t)
 Metallic behavior decreases going from left to right
across the Period table
● Increasing group number (left to right)
 Ability to lose electrons (form cations) becomes
more difficult with as Ionization Energy (IE)
increases
 Ability to gain electrons (form anions) increases
as Electron Affinity (EA) decreases (becomes
more negative)
● Elements on the left (more metallic) tend to form
positively charged “Cations”
● Elements on the right (more non-metallic) tend to
form negatively charged “Anions”
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Atomic Structure / Chemical Reactivity

Metallic Behavior (Con’t)
 Sodium (Na) group 1 – Very Metallic
● Readily loses electron (Na+ ion) which reacts
immediately with oxygen to form an oxide
 Aluminum (Al) group 3 – Metalloid
● Form some Al3+ ionic compounds, but is covalently
bonded in others
 Silicon (Si) group 4 – Metalliod
● Does not occur as a monoatomic ion
 Phosphorus (P) group 5 – non-metal
● Forms a few 3- ions
 Sulfur (S) group 6 – non-metal
● Forms 2- anions, such as Sulfide
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79
Atomic Structure / Chemical Reactivity

Metallic Behavior (Con’t)
 Acid-Base Behavior of Element Oxides
● Metals
 Most main group metals transfer electrons to
oxygen forming ionic oxides
 Ionic oxides act as bases producing
OH- (hydroxide) ions from O2● Non-metals
 Share electrons with oxygen to form covalent
oxides
 Covalent oxides act as acids producing
H+ ions (protons)
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Atomic Structure / Chemical Reactivity

Metallic Behavior (Con’t)
 Amphoteric Behavior
● Some metals and many metalloids form oxides
that can act as either an acid or a base
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Atomic Structure / Chemical Reactivity

Acid-Base behavior of common oxides
 As elements become more metallic going down a group,
the oxides become more basic
● Nitrogen Pentoxide (N2O5)
 Period 2 non-metallic forms nitric acid, a strong
acid
N2O5(s) + H2O(l) 2HNO3(aq)
 Tetraphosphorus decaoxide (P4O10)
 Slightly more metallic Period 3 non-metal forms a
weaker acid
P4O10(s) + 6H2O  4H3PO4(aq)
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Atomic Structure / Chemical Reactivity

Acid-Base behavior of common oxides
 Arsenic Pentoxide (As2O5)
● Group 4 metalloid (more metallic) is weakly basic
 Bismuth Pentoxide (Bi2O5)
● Group 5 metalloid (most metallic in group)
● Basic oxide, insoluble in water, forms salt & water
with an acid
Bi2O3(s) + 6HNO3(aq)  2Bi(NO3)3(aq) + 3H2O(l)
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Atomic Structure / Chemical Reactivity

Acid-Base behavior of common oxides
 Across a group
● Elements become less metallic across a group
● Oxides becomes more acidic
● Metallic Sodium (Na) (group 1) & Magnesium
(Mg) (group 2) form strongly basic oxides
● Metallic Aluminum (group 3) forms amphoteric
aluminum oxide (Al2O3), which can act as a base
to react with an acid or as an acid to react with a
base
Al2O3(s) + 6HCl(aq)  2AlCl3(aq) + 3H2O(l)
Al2O3(s) + 2NaOH(aq) + 3H2O(l) 
2NaAl(OH)4(aq)
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84
Atomic Structure / Chemical Reactivity

Acid-Base behavior of common oxides
 Across a group (Con’t)
● Silicon Dioxide (SiO4) group 4
 Weakly acidic forming salt & water with a base
SiO2(s) + 2 NaOH(aq)  Na2SiO3(aq) + H2O(l)
● Common oxides of Phosphous (group 5) Sulfur
(group 6) and Chlorine (group 7) are increasingly
acidic forming increasingly stronger acids
Acidity H3PO4 < H2SO4 < HClO4
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Atomic Structure / Chemical Reactivity

Acid-Base behavior of common oxides
 Trends in acid-base behavior of Group 5 and Period 3
oxides
● Red – Acidic (non-metal oxides)
● Blue – Basic (metal oxides)
● Other – Metalloid oxides (note gradations )
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Atomic Structure / Chemical Reactivity

Properties of Monoatomic Ions
 Electron Configuration of Main-Group ions
● Recall: Elements in Groups 1 & 2 readily lose
electrons to form cations and elements in groups 6 &
7 readily gain electrons to form anions
● The formation of the anions or cations result in a
filled outer shell, i.e., the nearest noble gas
configuration
Na(1s22s22p63s1)  Na+ (1s22s22p6)  Ne + eBr ([Ar] 4s23d104p5) + e-  Br- ([Ar] 4s23d104p6)
([Ar] 4s23d104p6)  [Kr]
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87
Atomic Structure / Chemical Reactivity

Properties of Monoatomic Ions (Con’t)
 Electron Configuration of Main-Group ions
● Energy to remove the outer valence shell electrons
(Ionization Energy) is supplied during the exothermic
reaction of a metal with a non-metal
● Removing more than one electron from Na to form
Na2+ or two electrons from Mg to form Mg3+ means
removing core (non-valence) electrons, which
requires much more energy than is available from the
chemical reaction
● Similarly, adding 2 electrons to Fluorine to form F2means adding electrons to the next energy level,
which would require a large amount of energy to
overcome the shielding of the nuclear charge by the
18 inner core electrons
Thus, compounds such as Na2F & Mg3O2 do not
exist
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88
Atomic Structure / Chemical Reactivity

Properties of Monoatomic Ions
 Electron Configuration of Main-Group ions
● Larger Metals of Groups 3, 4, 5
 Energetically impossible for them to lose enough
electrons to attain noble gas configuration
Tin (Sn) [Kr] 5s25p24d10 would have to lose 14
electrons (two 5p, ten 4d, and two 5s) to be
isoelectronic with Krypton (Kr) – [Ar] 4s24p6
 Cations formed through a different process
 Sn+4 – loss of two 5s & two 5p electrons,
attaining stability from the filled in 4d sublevel
 Sn+2 – loss of two 5p electrons, attaining
stability from the filled 5s & 4d sublevels
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89
Atomic Structure / Chemical Reactivity

Properties of Monoatomic Ions
 Electron Configuration of Main-Group ions
● Larger Metals of Groups 3, 4, 5
 Carbon
 Would have to either lose 4 electrons to attain
the C4+ Helium configuration or gain 4
electrons to attain the C4- Neon configuration
 In either case, the energy requirements are
extremely high, i.e. sun-like temperatures of
106k
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90
Atomic Structure / Chemical Reactivity

Properties of Monoatomic Ions
 Electron Configuration of Main-Group ions
● Most elements that form Monatomic ions that are
Isoelectronic with a noble gas lie in the four groups
that flank group 8
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91
Practice Problem
Using condensed electron configurations, write reactions for
the formation of the common ions of the following:
Iodine:
I ([Kr] 5s24d105p5) + e-  I- ([Kr] 5s24d105p6)  I- [Xe]
Potassium:
K ([Ar] 4s1)  K+ ([Ar]) + eIndium:
Group 3A – loses 3 electrons or loses 1 electron
In ([Kr] 5s24d105p1)  In3+ ([Kr] 4d10)
+ 3e-
In ([Kr] 5s24d105p1)  In+ ([Kr] 5s24d10) + 1e1/13/2015
92
Atomic Structure / Chemical Reactivity

Electron Configurations of Transition Metal Ions
 Transition metal ions rarely attain noble gas
configurations
 Energy required to attain noble gas configuration is
very high
● Exceptions
 Scandium – forms Sc3+; Titanium – forms
Ti4+
 In Periods 4 & 5, a transition metal can form
more than one cation by losing all of its ns
and some of the (n-1)d electrons
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93
Atomic Structure / Chemical Reactivity
 Electron Configurations of Transition Metal Ions
 Aufbau electron build-up
 At the beginning of Period 4, the 4s orbital is nearer
the nucleus than the 3d orbital making it more stable
than the empty 3d orbital
 The first & second electrons fill the 4s orbital before
filling the empty 3d orbitals
● At the beginning of the transition elements (group
3B), however, the previously filled 4s orbitals do not
do a very good job of shielding the 3d electrons
● The 3d orbitals now become more stable than the 4s
orbitals and begin to fill under the influence of
increased nuclear charge - a cross-over in orbital
energy
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94
Atomic Structure / Chemical Reactivity
 Electron Configurations of Transition Metal Ions
 Aufbau electron build-up (Con’t)
 The 4s electrons, which were added before the 3d
electrons, are now lost preferentially before the 3d
electrons to form the transition metal electrons
 Simple Rules for forming the ion of any “Main Group” or
“Transition” Group element
● Electrons with the highest “n” value are removed first
● For main-group, s block metals, remove all electrons
with the highest “n” value
● For main-group, p-block metals, remove “np”
electrons before “ns” electrons
● For transition (d-block) metals, remove “ns” electrons
before “(n-1)d” electrons
● For non-metals, add electrons to the “p” orbitals of
the highest “n” value
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Magnetic Properties

A spinning electron behaves like a tiny magnet generating a
magnetic field
 A single electron (unpaired) in an orbital can be affected by
an externally applied magnetic field
 A Paramagnetic element (or ion) has 1 or more orbitals
with unpaired electrons and is weakly attracted by a
magnetic field
Titanium
[Ar]4s23d2
4s
3d
4p
4s
3d
4p
 A Diamagnetic element (or ion) has only paired electrons
and is not attracted by a magnetic field
Copper ion Cu+ [Ar]4s23d10
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96
Atomic Structure / Chemical Reactivity

Magnetic Properties of Transition Metal Ions
 Ag (Z=47) [Kr] 5s14d10
5s
4d
5p
Unpaired – Paramagnetic – split by applied magnetic field
Cd (Z=48) [Kr] 5s24d10
5s
Paired
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4d
5p
– Diamagnetic – not split by applied magnetic
field
97
Atomic Structure / Chemical Reactivity

Using Paramagnetism to verify electron configuration
 Titanium
(Ti)
[Ar] 4s23d2
 Titanium (II) Ion (Ti2+) ([Ar] 3d2) + 2eTi
4s
3d
4p
Ti2+
4s
3d
4p
Paramagnetic
Unpaired e-
Paramagnetic
Unpaired e-
If Titanium had lost its two “3d” electrons, the Titanium Ion would have
been “diamagnetic (all electrons shared)
The Titanium Ion actually shows properties of “Paramagnetism”
(The mass of the titanium ion is affected when placed in a magnetic field)
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Atomic Structure / Chemical Reactivity

Increasing “Paramagnetism”
 Iron (Fe)  Iron III (Fe3+)
Fe ([Ar] 4s23d6)  Fe3+ ([Ar] 3d5) + 3eFe
4s
3d
4p
4s
3d
4p
Fe3+
The loss of the 2 4s electrons and
one of its paired 3d electrons
results in “increased” paramagnetism
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99
Practice Problem
Use condensed electron configuration to write the reaction
for the formation of Mn2+ ion, and predict whether the ion is
paramagnetic
Manganese Mn (Z = 25)
Mn ([Ar] 4s23d5) → Mn2+ ([Ar] 3d5) + 2eMn
4s
3d
4p
4s
3d
4p
Mn2+
Rule: Remove “ns” electrons first
The Mn2+ ion is “paramagnetic”
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100
Practice Problem
Use condensed electron configuration to write the reaction
for the formation of Cr3+ ion, and predict whether the ion is
paramagnetic
Chromium - Cr (Z = 24)
Cr ([Ar] 4s13d5) → Cr3+ ([Ar] 3d3) + 3eCr
4s
3d
4p
4s
3d
4p
Cr3+
Note irregularity for Cr: 4s subshell fills before 3d subshell is complete
Rule: Remove “ns” electrons first
The Cr3+ ion is “paramagnetic”
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101
Summary Equations
N = Principal quantum number (size, energy) values = 1, 2, 3 ….
l
= Angular Momentum (orbital shape)
0(s), 1(p), 2(d), 3(f)
values = 0  n-1
ml = magnetic (orbital orientation)
values = -l ... 0 ... +l
Ms = Spin (direction)
values
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= -1/2 & +1/2
102