Chapter Seven:
Atomic Structure and Periodicity
Arrangement of Electrons
in Atoms
Reminder: electrons in atoms are arranged as…
SHELLS (n)
SUBSHELLS (l)
ORBITALS (ml)
Electrons in Atoms
When n = 1, then l = 0
This shell has a single orbital (1s) to which 2
electrons can be assigned.
When n = 2, then l = 0, 1
2s orbital
three 2p orbitals
TOTAL =
2 electrons
6 electrons
8 electrons
Electrons in
Shells
The number of
shells increases
with the shell value
(n).
Therefore each
higher shell holds
more electrons.
(more orbitals)
Quantum Numbers
Electrons in an atom are arranged by:
(quantum number)
Principle energy levels (shells)
n
angular energy levels (sub-shells)
l
oriented energy levels (orbitals)
ml
Within an individual orbital, there may only be two electrons
differentiated by their spin.
Spin states
(electrons)
ms
Atomic Subshell Energies
& Electron Assignments
Based on theoretical and experimental studies of
electron distributions in atoms, chemists have found
there are two general rules that help predict these
arrangements:
1.
Electrons are assigned to subshells in order of
increasing “n + l” value.
2. For two subshells with the same value of “n + l”
electrons are assigned first to the subshell of lower n.
Single-Electron Atom Energy Levels
Energy
4s
3s
3p
2s
2p
3d
In a Hydrogen atom (1–electron)
the orbitals of a subshell are
equal in energy (degenerate)
1s
Multi-Electron Atom Energy Levels
3d
4s
3p
E=n+l
3s
2p
2s
For multi-electron atoms,
screening results in the 4sorbital having a lower energy
that that of the 3d-orbital.
4s: n + l = 4 + 0 = 4
vs.
3d: n + l = 3 + 2 = 5
1s
Effective Nuclear Charge, Z*
Z* is the net charge experienced by a particular
electron in a multi-electron atom resulting from a
balance of the attractive force of the nucleus and the
repulsive forces of other electrons.
 Z* increases across a period owing to incomplete
screening by inner electrons
 This explains why E(4s electron) < E(3p electron)
 Z*  [ Z  (no. inner electrons) ]

Charge felt by 2s electron in: Li Z* = 3  2 = 1
Be Z* = 4  2 = 2
B
Z* = 5  2 = 3
and so on!
Electron Configurations &
the Periodic Table
Electron Filling Order
Assigning Electrons to
 In the H-atom, all
Subshells

1.
2.
subshells of same n have
same energy.
In a multi-electron atom:
subshells increase in
energy as value of n + l
increases.
for subshells of same n + I,
the subshell with lower n
is lower in energy.
Orbital Filling Rules:
Aufbau Principle: Lower energy orbitals fill first.
Hund’s Rule: Degenerate orbitals (those of the same energy)
are filled with electrons until all are half filled before pairing up
of electrons can occur.
Pauli exclusion principle: Individual orbitals only hold two
electrons, and each should have different spin.
“s” orbitals can hold 2 electrons
“p” orbitals hold up to 6 electrons
“d” orbitals can hold up to 10 electrons
Orbital Filling: The “Hund’s Rule”
Hund’s Rule.
Degenerate orbitals are filled with electrons
until all are half-filled before pairing up of
electrons can occur.
Consider a set of 2p orbitals:



2p
Electrons fill in this manner
Orbital Filling: The “Pauli Principle”
“Pauli exclusion principle”
Individual orbitals only hold two electrons, and
each should have opposite spin.
Consider a set of 2p orbitals:



2p
*the convention
is to write up,
then down.



 = spin up
 = spin down
Electrons fill in this manner
Electron Configuration: Orbital Box
Notation
Electrons fill the orbitals from
lowest to highest energy.
The electron configuration of an
atom is the total sum of the
electrons from lowest to highest
shell.
Example:
Nitrogen: N has an atomic number of 7, therefore 7 electrons
Orbitals

1s



2s



2p
Electron Configuration (spdf) notation:
1s2 2s2 2p3
Atomic Electron Configurations
Carbon
Group 4A
Atomic number = 6
6 total electrons
1s2 2s2 2p2
3p
3s
2p
2s

1s
1s



2s


2p
Electron Configuration in the
3rd period:
Orbital box notation:
1s
2s
2p
3s
3p
3p
This corresponds to the energy
level diagram:
3s
2p
2s
1s
Electron Configuration in the 3rd
Period
Orbital box notation:


  

1s
2s
2p
3s

3p

3p

Aluminum: Al (13 electrons)
spdf Electron Configuration
3s

2s
1s2 2s2 2p6 3s2 3p1

1s


2p

Noble Gas Notation
The electron configuration of an element can be represented as a
function of the core electrons in terms of a noble gas and the
valence electrons.
Full electron configuration spdf notation
1s22s22p63s23p2
Orbital Box Notation


[Ne]

3s

3p
Noble gas Notation
[Ne] 3s23p2
Noble Gas Notation

The innermost electrons (core) can be represented by the full shell
of noble gas electron configuration:
1s22s2 = [He], 1s22s22p6 = [Ne], 1s22s22p63s23p6 =
[Ar]...
Element

Full Electron Config.
Noble Gas Notation
The outermost electrons are referred to as the “Valence”
electrons”.
Mg
1s2 2s2 2p6 3s2
[Ne] 3s2
Transition Metal
All 4th period and beyond d-block elements
have the electron configuration [Ar] nsx (n - 1)dy
Where n is the period and x, y are particular to
the element.
Chromium
Iron
Copper
Transition Elements
Electron Configurations are written by shell even
though the electrons fill by the periodic table:
Ni:
last electron to fill: 3d8
electron configuration by filling:
1s2 2s2 2p6 3s2 3p6 4s2 3d8
electron configuration by shell:
(write this way)
1s2 2s2 2p6 3s2 3p6 3d8 4s2
Lanthanides & Actinides
f-block elements: These elements have the
configuration [core] nsx (n - 1)dy (n - 2)fz
Where n is the period and x, y & z are particular
to the element.
Cerium:
[Xe] 6s2 5d1 4f1
Uranium:
[Rn] 7s2 6d1 5f3
Some Anomalies
Some
irregularities
occur when
there are
enough
electrons to
half-fill s and d
orbitals on a
given row.
Some Anomalies
For instance,
the electron
configuration
for copper is
[Ar] 4s1 3d10
rather than the
expected
[Ar] 4s2 3d9.
Some Anomalies

This occurs
because the 4s
and 3d orbitals
are very close in
energy.

These
anomalies
occur in f-block
atoms, as well.
Filling Rules

Aufbau Principle:
 Electrons
first.

fill the lowest energy levels
Pauli Exclusion Principle:
 Electrons
can fill two/orbital, providing
they have opposite spins ( )

Hund’s Rule:
 Orbitals
of the same energy level
(p,d,f) distribute electrons 1/orbital
before adding the second.
Who are they?
1.
1S22S22P63S2
2.
[KR]5S24D105P5
3.
1S22S22P3
4.
[RN]7S1
5.
[AR]4S23D10
Who are they?
1.
1s22s22p63s2 Magnesium
2.
[Kr]5s24d105p5 Iodine
3.
1s22s22p3 Nitrogen
4.
[Rn]7s1 Francium
5.
[Ar]4s23d10 Zinc
What’s wrong?
Mg: [Ar]3s2
Fe: 1s22s22p63s23p64s24d6
Al: [Ne] 3s3
Mistakes!
Mg: [Ar]3s2
Fe: 1s22s22p63s23p64s24d6
Al: [Ne] 3s3
Corrections
Mg: [Ne]3s2
Fe: 1s22s22p63s23p64s23d6
Al: [Ne] 3s23p1
Valence Electrons:

Electrons at the highest energy level.

Electrons available to be gained/lost/shared in a chemical
reaction

Valence electron configuration: nsxpx (total of 8 valence electrons)

Representation of valence electrons in Lewis dot notation:
Formation of ions &
Valence
electrons
Ions are formed when atoms either:

1. Cation (+): Give up (lose) electrons
2.
Anion (-): Gain electrons
•
Overall goal: stable noble gas notation
•
Elements with < 4 valence electrons- form cations
•
Elements with > 4 valence electrons for anions.
•
Non-metals with 4 valence electrons- do not form ions
•
Noble gases (8 valence electrons) – are unreactive
Ions & Electron
Configuration

Atoms or groups of atoms that carry a charge

Cations- positive charge




Formed when an atom loses electron(s)
Na  Na+ + eNa+ : 1s22s22p6 (10 e-) = [Ne]
Anions –negative charge



Formed when an atom gain electrons(s)
F + e-  FF- : 1s22s22p6 (10 e-) = [Ne]
Isoelectronic series:

Ions that contain the same number of
electrons.

Example: Al3+, Mg 2+, Na+, F-, O2-, N3-
Transition Metals & Ions

Transition metals: multi-valent ions
(more than one possible charge)

Electrons are removed from the highest quantum
number first:

Example: Cu+ & Cu 2+
Cu+: [Ar] 3d10
Cu2+ : [Ar] 3d9
Electron Spin &
Magnetism • Diamagnetic Substances:
Are NOT attracted to a
magnetic field
• Paramagnetic Substances:
ARE attracted to a magnetic
field.
• Substances with unpaired
electrons are paramagnetic.
Electron Configurations of
Ions
Ions with UNPAIRED ELECTRONS are
PARAMAGNETIC (attracted to a magnetic
field).
Ions without UNPAIRED ELECTRONS are
DIAMAGNETIC (not attracted to a magnetic
field).
Fe3+ ions in Fe2O3 have
5 unpaired electrons.
This makes the sample
paramagnetic.
Practice: Electron
Configuration

Write the following electron
configurations
1.
Silicon- orbital notation
2.
Strontium – long (spdf) notation
3.
Bismuth – noble gas notation
4.
Silver – noble gas notation
5.
O 2- - orbital notation
6.
Fe2+ - noble gas notation
Practice #2: Electron
Configuration

Write the following electron
configurations
1.
Chromium- long (spdf) notation
2.
Sulfur – orbital notation
3.
Ag+ – noble gas notation
4.
Americium – noble gas notation
5.
Mg 2+ - orbital notation
6.
Krypton – noble gas notation
Development of Periodic Table

Elements in the
same group
generally have
similar chemical
properties.

Physical properties
are not necessarily
similar, however.
Periodic Trends

In this chapter, we will rationalize
observed trends in:
 Sizes
of atoms and ions.
 Ionization
 Electron
energy.
affinity.
General Periodic Trends



Atomic and ionic size
Ionization energy
Electron affinity
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.
Effective Nuclear Charge
In a many-electron
atom, electrons are
both attracted to the
nucleus and repelled
by other electrons.
 The nuclear charge
that an electron
experiences
depends on both
factors.

Effective Nuclear Charge
The effective nuclear
charge, Zeff, is found
this way:
Zeff = Z − S
where Z is the atomic
number and S is a
screening constant,
usually close to the
number of inner
electrons.
What Is the Size of an Atom?
The bonding
atomic radius is
defined as onehalf of the
distance between
covalently
bonded nuclei.
Sizes of Atoms
Bonding atomic radius tends
to…
…decrease from left to right
across a row
(due to increasing Zeff).
…increase from top to bottom
of a column
(due to increasing value of n).
Ionization Energy

The ionization energy is the amount of
energy required to remove an electron
from the ground state of a gaseous atom
or ion.
 The
first ionization energy is that energy
required to remove first electron.
A(g) + energy  A+ + e-

 The
second ionization energy is that energy
required to remove second electron, etc.

A+(g) + energy  A2+ + e-
Ionization Energy
It requires more energy to remove each
successive electron.
 When all valence electrons have been
removed, the ionization energy takes a
quantum leap.

Trends in First Ionization Energies

As one goes down
a column, less
energy is required
to remove the first
electron.
 For
atoms in the
same group, Zeff is
essentially the same,
but the valence
electrons are farther
from the nucleus.
Trends in First Ionization Energies

Generally, as one
goes across a row,
it gets harder to
remove an
electron.
 As
you go from left
to right, Zeff
increases.
Trends in First Ionization Energies
However, there
are two apparent
discontinuities in
this trend.
Trends in First Ionization Energies


The first occurs
between Groups IIA
and IIIA.
In this case the
electron is removed
from a p-orbital rather
than an s-orbital.

The electron removed is
farther from nucleus.

There is also a small
amount of repulsion by
the s electrons.
Electron Affinity ( EAH)
Electron affinity is the energy
change accompanying the
addition of an electron to a
gaseous atom:
Cl(g) + e−  Cl−
Trends in Electron Affinity
In general,
electron affinity
becomes more
exothermic as you
go from left to
right across a row.
Trends in Electron Affinity
There are
again,
however, two
discontinuities
in this trend.
Trends in Electron Affinity

The first occurs
between Groups 1
and 2.
 The
added electron
must go in a p-orbital,
not an s-orbital.
 The
electron is farther
from nucleus and
feels repulsion from
the s-electrons.
Trends in Electron Affinity

The second occurs
between Groups
IVA and VA.
 Group
VA has no
empty orbitals.
 The extra electron
must go into an
already occupied
orbital, creating
repulsion.
Ions:

Write the noble gas notation for the
following ions:
 Mg2+
 P3 Fe3+

Write the noble gas notation for
the following ions:
 Mg2+
[He]2s22p6
 P3-
[Ne] 3s23p6
 Fe3+
[Ar] 3d5
Sizes of Ions

Ionic size
depends upon:
 The
nuclear
charge.
 The
number of
electrons.
 The
orbitals in
which electrons
reside.
Sizes of Ions

Cations are
smaller than their
parent atoms.
 The
outermost
electron is
removed and
repulsions
between
electrons are
reduced.
Sizes of Ions

Anions are larger
than their parent
atoms.
 Electrons
are
added and
repulsions
between
electrons are
increased.
Sizes of Ions

Ions increase in
size as you go
down a column.
 This
is due to
increasing value of
n.
Sizes of Ions

In an isoelectronic series, ions have the
same number of electrons.

Ionic size decreases with an increasing
nuclear charge.
Problem:
Rank the following ions in order of decreasing size?
Na+, N3-,
Mg2+, F–
O2–
Problem:
Rank the following ions in order of decreasing size?
Na+, N3-,
Mg2+, F–
O2–
Ion
Na+
N3Mg2+
F–
O2–
# of protons
# of electrons
ratio of e/p
Problem:
Rank the following ions in order of decreasing size?
Na+, N3-,
Mg2+, F–
O2–
Ion
# of protons
Na+
11
N3-
7
Mg2+
12
F–
9
O2–
8
# of electrons
ratio of e/p
Problem:
Rank the following ions in order of decreasing size?
Na+, N3-,
Mg2+, F–
O2–
Ion
# of protons
# of electrons
Na+
11
10
N3-
7
10
Mg2+
12
10
F–
9
10
O2–
8
10
ratio of e/p
Problem:
Rank the following ions in order of decreasing size?
Na+, N3-,
Mg2+, F–
O2–
Ion
# of protons
# of electrons
ratio of e/p
Na+
11
10
0.909
N3-
7
10
1.43
Mg2+
12
10
0.833
F–
9
10
1.11
O2–
8
10
1.25
Ion
e/p ratio
Na+
0.909
N3-
1.43
Mg2+
0.833
F–
1.11
O2–
1.25
Since N3- has the highest ratio of
electrons to protons, it must have the
largest radius.
Ion
e/p ratio
Na+
0.909
N3-
1.43
Mg2+
0.833
Since N3- has the highest ratio of
electrons to protons, it must have the
largest radius.
Since Mg2+ has the lowest, it must have
the smallest radius.
The rest can be ranked by ratio.
F–
1.11
O2–
1.25
Ion
e/p ratio
Na+
0.909
N3-
1.43
Mg2+
0.833
Since N3- has the highest ratio of
electrons to protons, it must have the
largest radius.
Since Mg2+ has the lowest, it must have
the smallest radius.
The rest can be ranked by ratio.
F–
1.11
O2–
1.25
N3-
>
O2– >
F– >
Na+ >
Mg2+
Decreasing size
Notice that they all have 10 electrons: They are
isoelectronic (same electron configuration) as
Ne.
Summary of Periodic
Trends
Moving through
the periodic
table:
Down a group
Across a Period
Atomic
radii
Increase
Decrease
Ionization
Energy
Electron
Affinity
Decrease
Becomes
less
exothermic
Increase
Becomes
more
exothermic