Chapter 7
The Structure of Atoms and
Periodic Trends
Arrangement of Electrons in Atoms
Electrons in atoms are arranged as:
Shells (n)
Subshells (l)
Subshell orientation (ml)
Pauli’s Exclusion Principle
discovered in 1925 by Wolfgang Pauli
-No two electrons in an atom can have the same
set of 4 quantum numbers
Practice: What are the 4 quantum numbers for
each electron in He?
Aufbau Principle
Describes the electron filling order in atoms
-electrons are placed in the lowest available energy
orbital
-the periodic table is a
function of electron
configurations for the
elements
Electron Configuration
To remember the correct filling order for
electrons in atoms:
Electron Configuration
Writing Electron Configurations
Two ways to express electron configuration:
1. spdf notation
Example: H atomic number = 1
1s
1
no. of
electrons
value of l
value of n
Writing Electron Configurations
2. Orbital box notation
spdf notation
ORBITAL BOX NOTATION
for He, atomic number = 2
Arrows
2
depict
electron
spin
1s
1s
Electron Configurations
Using the Aufbau Principle to determine the
electronic configurations of the elements
1st row elements:
1s

1
H
2
He 
Configurat ion
1
1s
1s
2
Electron Configurations
Hund’s rule: electrons
fill suborbitals by
placing electrons in
each suborbital
unpaired first with the
same spin direction,
then the electrons pair
Electron Configurations
3s
3p
Configurat ion
11 Na Ne 
12
Mg Ne 
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
Ne
Ne
Ne
Ne
Ne
Ne



 

  

  

  

  
Ne 3s1
Ne 3s 2
Ne 3s 2 3p1
Electron Configurations and Quantum
Numbers
We can write a complete set of quantum
numbers for all of the electrons in every
element:
– Na
– Ca
– Fe
3s
11
Na Ne 
3p
Configurat ion
Ne 3s
1
Electron Configurations and Quantum
Numbers
n
l
ml
1
0
0
2 nd e - 1
0
0
3rd e -
2
0
0
4 th e -
2
0
0
5th e -
2
1
6 th e -
2
1
7 th e -
2
1
8th e -
2
1
9 th e -
2
1
10 th e -
2
1
11th e -
3
0
1st e -
ms
 1/2 
1 s electrons
 1/2 
 1/2 
2 s electrons
 1/2 
 1/2 

0
 1/2 
 1  1/2 

2 p electrons
 1  1/2 
The ml and ms are interchangeable
0
 1/2 

 1  1/2 

0
 1/23 s electron
-1
Electron Configurations and Quantum
Numbers
Noble Gas Notation (or short hand notation):
The first 18 electrons in Ca are represented with the
preceding noble gas ([Ar])
- we only concern ourselves with the outermost e-
3d
20
Ca [Ar]
4s

4p
Configurat ion
Ar  4s
2
Skip the first 18 electrons
Electron Configurations and Quantum
Numbers
n
l
ml
[Ar ]19 th e -
4
0
0
20 th e -
4
0
0
ms
 1/2 

4 s electrons
 1/2 

Electron Configurations and Quantum
Numbers
There is only one set of 4 quantum numbers for each of
the 26 electrons in Fe:
– To save space, we use the symbol [Ar] to represent the first
18 electrons in Fe
3d
26 Fe Ar      
4s

4p
Configurat ion
Ar  4s2 3d6
Electron Configurations of Ions
Electrons are removed from subshell of
highest energy level (n-level)
P0 [Ne] 3s2 3p3
-3e- ---> P3+ [Ne] 3s2 3p0
3p
3p
3s
3s
2p
2p
2s
2s
1s
1s
Electron Configurations of Ions
For transition metals, remove the highest
s-orbital electrons first:
Fe [Ar] 4s2 3d6
-2 electrons
Fe2+ [Ar] 3d6
-3 electrons
Fe3+ [Ar] 3d5
To form cations, always remove electrons of highest n value first!
More About the Periodic Table
Representative Elements
Groups IA, IIA, IIIA-VIIIA
– These elements will have
their “outermost” electron
in an outer s or p orbital
– Variations in their properties
are similar from top-tobottom
More About the Periodic Table
d-Transition Elements
All have d electrons
-With n s-orbitals
-With n-1 d–orbitals
Have small property
variations from row-torow
More About the Periodic Table
f - transition metals
-Sometimes called inner
transition metals
-Electrons are being
added to f orbitals
Extremely small variations in
properties from one element
to another
More About the Periodic Table
Noble Gases
-Have filled electron shells
-have similar chemical reactivities
-similar electronic structures
He
Ne
Ar
Kr
Xe
Rn
1s2
[He] 2s2 2p6
[Ne] 3s2 3p6
[Ar] 4s2 4p6
[Kr] 5s2 5p6
[Xe] 6s2 6p6
Periodic Properties
• Atomic radii describes the relative sizes of atoms
• Atomic radii increase within
a column
• Atomic radii decrease within
a row
Periodic Properties
Example: Arrange these elements based on their
atomic radii:
Se, S, O, Te
O < S < Se < Te
Periodic Properties
Example: Arrange these elements based on their
atomic radii:
P, Cl, S, Si
Cl < S < P < Si
Periodic Properties
Electronegativity: measure of the tendency of an
atom to attract electrons to itself
-Fluorine is the most electronegative element
-Cesium is the least electronegative element
Electronegativity increase from left-to-right and
decrease from top-to-bottom
increase
decrease
Periodic Properties
Example: Arrange these elements based on their
electronegativity:
Se, Ge, Br, As
Ge < As < Se < Br
Periodic Properties
Example: Arrange these elements based on their
electronegativity:
Be, Mg, Ca, Ba
Ba < Ca < Mg < Be
Periodic Properties
Ionization Energy: energy required to remove
an electron from an atom in the gas state
First ionization energy (IE1)
– Energy required to remove the first electron from an
atom in the gas state to form a 1+ ion
Atom(g) + energy  Atom+(g) + eExample:
Mg(g) + 738kJ/mol  Mg+ + e-
Periodic Properties
Second ionization energy (IE2)
– The amount of energy required to remove
the second electron from a gaseous 1+ ion
Atom+ + energy  Atom2+ + eMg+ + 1451 kJ/mol  Mg2+ + e- Atoms can have 3rd (IE3), 4th (IE4), etc.
- Each IE is significantly higher than the
previous IE
Periodic Properties
Ionization Energy:
1. IE2 > IE1
always takes more energy to remove a second
electron from an ion
2. IE1 increases to the right
Important exceptions are Be & Mg, N & P, etc. due to
filled and half-filled subshells
3. IE1 decrease down
First Ionization Energies
He
2500
Ne
2000
Ionization
Energy
(kJ/mol)
N
1500
1000
H
C
Be
F
Ar
Cl
P
O
Mg
S
B
500
Li
Ca
Si
Na
Al
K
0
1 2
3 4
5 6 7
8 9 10 11 12 13 14 15 16 17 18 19 20
Atomic Number
Periodic Properties
Example: Arrange these elements based on their
first ionization energies:
Sr, Be, Ca, Mg
Sr < Ca < Mg < Be
Periodic Properties
Example: Arrange these elements based on their
first ionization energies:
Al, Cl, Na, P
Na < Al < P < Cl
Periodic Properties
Electron Affinity: Energy absorbed when an
electron is added to an atom to form a negative
ion
Sign conventions for electron affinity:
– If electron affinity > 0 energy is absorbed
– If electron affinity < 0 energy is released
Electron affinity is the measure of an atom’s
ability to form negative ions
atom(g) + e- + EA 
atom-(g)
Periodic Properties
Examples of electron affinity values:
Mg(g) + e- + 231 kJ/mol  Mg-(g)
EA = +231 kJ/mol
Br(g) + e-  Br-(g) + 323 kJ/mol
EA = -323 kJ/mol
Increasing ability to
decreasing ability
to add electrons
add electrons
Electron Affinity
Electron Affinities of Some Elements
Electron Affinity (kJ/mol)
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
0
-50
-100
-150
-200
-250
-300
-350
-400
He
Be
B
N
Ne
Mg
Al
Ar
P
Na
H
Li
K
O
C
Si
S
F
Atomic Number
Ca
Cl
Periodic Properties
Example: Arrange these elements based on their
electron affinities:
Al, Mg, Si, Na
Si < Al < Na < Mg
Periodic Properties
Ionic Radius: diameter of an atom in its ionized form
-Cations are always smaller
Element
Atomic
Radius (Å)
Ion
Ionic
Radius (Å)
Li
Be
1.52
1.12
Li+
Be2+
0.90
0.59
Periodic Properties
Anions are always larger
Element
N
O
F
Atomic
Radius(Å)
Ion
0.75
0.73
0.72
N3-
O2-
F1-
Ionic
Radius(Å)
1.71
1.26
1.19
Periodic Properties
Cation radii decrease from left to right across a
period
– Increasing nuclear charge attracts the electrons and
decreases the radius.
Ion
Rb+
Sr2+
In3+
Ionic
Radii(Å)
1.66
1.32
0.94
Periodic Properties
Anion radii decrease from left to right across a
period
– Increasing electron numbers in highly charged ions
cause the electrons to repel and increase the ionic
radius
Ion
N3-
O2-
F1-
Ionic
Radii(Å)
1.71
1.26
1.19
Ionic Radii
Active Figure 8.15
Periodic Properties
Example: Arrange these elements based on their
ionic radii:
Ca2+, K+, Ga3+
K1+ > Ca2+ > Ga3+
Periodic Properties
Example: Arrange these elements based on their
ionic radii:
Cl-1, Se-2, Br-1, S-2
Cl1- < S2- < Br1- < Se2-