Electron Configuration and
Atomic Properties
Topics:
Electron Spins and Magnetism
Orbital Energy
Electronic Configurations of Elements
Atomic Properties
Ions
OWLs Due on 28-November.
Electron Spin and Magnetism

There is an
additional quantum
number, and one
that is very
important!
◦ Spin Quantum
Number (ms)
◦ Electrons can have a
spin of +1/2 or -1/2
◦ Since electrons are
charged particles, as
they move (spin, for
example), they
create magnetic
fields.

We’ll see how we fill electrons with spins
into orbitals soon, but we need to know
that there are 3 types of magnetic materials:
◦ Diamagnetic (non-magnetic)
◦ Paramagnetic (weakly magnetic)
◦ Ferromagnetic (strongly magnetic)
Orbital Energies (single e- species)

In single electron
species (hydrogen) or
even some ions, the
orbitals at each energy
level have the same
energy.
◦ Even if the orbitals are
different in physical size

This is a very simplistic
model that works for
few species
Orbital Energies (multiple e- species)

In multiple electron
species, orbital energies
are intermixed among
different “n” energy levels.
◦ Most common
◦ Leads to electron
configurations and systematic
filling for most elements and
ions.

The energies change
because electrons now
interact with other
electrons
◦ Repulsive forces
Electron Configuration in Atoms

The Pauli Exclusion
Principle states, simply,
that no two electrons
may have the same set of
quantum numbers.
◦
◦
◦
◦

n
l
ml
ms
Additionally, atomic
orbitals are filled from
the lowest energy up
when the atom is in the
“ground” state
◦ Lowest energy state
Electron Configuration in Atoms

This results in the following order:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p,
6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p



You should know up through 4f
The total number of electrons is the same as the
atomic number
We have as many unpaired electrons in a specific
orbital (s, p, d, f) as we can fit.
◦ Hund’s Rule

And what does this mean for the Periodic Table?
◦ Flashback to the previous chapter……
Let’s Build Some Electron
Configurations!

Remember that there are always a few
exceptions
◦ Focus on the rules rather than the
exceptions.

Remember your filling order.
◦
◦
◦
◦
Fill in a box diagram
Fill in an energy level diagram
Write it in notation
Use the atomic number and position of the
element on the periodic table to help!!!!

Hydrogen (ground state):
◦
◦
◦
◦
◦
◦

Atomic Number 1
# of electrons = 1
n=1
l = 0 (s orbital)
ml = 0 (one s orbital)
ms = +1/2
1s1

Helium (ground state)  Lithium (ground state)
◦
◦
◦
◦
◦
◦

◦
◦
◦
◦
◦
◦
◦
Atomic Number 2
# of electrons = 2
n=1
l = 0 (s orbital)
ml = 0 (one s orbital)
ms = +1/2 and -1/2
1s2

Atomic Number 3
# of electrons = 3
n = 1, 2
l = 0,1 (s and p orbitals)
ml = 0 (one s orbital)
ml = -1, 0, 1 (3 p orbitals)
ms = +1/2 and -1/2
1s2 2s1

Carbon (ground state)
◦
◦
◦
◦
◦
◦
Atomic Number 6
# of electrons = 6
n = 1, 2 (second period on table)
l = 0, 1 (s and p orbitals)
ml = 0 (one s orbital), -1,0,1 (3 p orbitals)
ms = +1/2 and -1/2
◦ 1s22s22p2
What element is this?
What is the electronic
configuration for Calcium?
Shorthand Notation……

Sometimes we abbreviate electron
configurations using Noble Gas Notation:
◦ What element is this?
◦ [Ar]3d104s24p5
Scandium
(Sc)
Back to the Periodic Table…..
Mendeleev’s periodic table
There are other types of electron
configurations we need to consider
(and terms)

Ground (lowest) versus excited (higher)
energy state….

Inner (core) shell and valence (outer)
shell electrons

Atoms versus ions (a topic for later in
Chapter 7)
Ground vs. Excited State

An excited state electronic configuration is
present when an atom (or ion) absorbs
energy and an electron is promoted to a
higher energy level
◦ This can occur even if the n energy level does not
exist in the ground state. The level is still there.
◦ Ground State Mg (1s22s22p63s2), 12 electrons.
◦ Excited State Mg (just one of many possible)
 1s22s22p63s13p1
 The atom must absorb energy for this to happen.
 When it transitions back to the ground state, that exact
amount of energy is given off
 Light
 Heat
 Kinetic energy transferred to another atom or molecule.
Inner versus Outer Shell

Inner (core) shell electrons are those in full
(“closed”) n energy levels
◦ Ordinarily those seen in the noble gas
configuration.
◦ Take Aluminum for example
 1s22s22p63s23p1
 [Ne] 3s23p1
 [Ne] electrons represent the inner or core shell
electrons
 The 3s2 and 3p1 electrons are the valence or outer
shell.
 When ions are formed, only valence electrons are
gained or lost
 Al3+
Periodic Trends and Properties

Effective Nuclear Charge (Z*)
◦ As atomic number increases, so does the number
of protons. The nuclear charge increases, which
raises the energy of orbitals surrounding the
nucleus.
◦ Effective nuclear charge (Z*) = Z – (# of core
shell electrons)
◦ It is a relative number, used for basic
comparisons.
◦ It represents the nuclear charge experienced by
the highest energy, valence electrons
Atomic Size
Closely related to orbital configuration, energies
and effective nuclear charge
 Atoms with a greater nuclear charge “pull”
electrons in closer to the nucleus and are smaller

◦ Opposites attract
Covalent radius: Distance between two nuclei when two atoms are
bonded together (Cl2 as an example)
Metallic radius: Distance between nuclei when atoms of a metallic
element are near each other in a metallic crystal (say a block of Zn)
Ionization Energy

The energy required to remove an
electron from an element in a gaseous
state.

Increases across a period because of
increasing orbital energy and effective
nuclear charge
Ionization energy trends
Ionization energy trends
Electron Affinity

The energy change when a gaseous atom
gains an electron

More negative values represent a greater
affinity.
Ions…

Filled shells (n energy levels) or orbitals,
represent the most stable electron
configurations
◦ Electrons of opposite spin are paired
◦ Energy levels may be full
Ions form because the energy level of that
electron configuration is particularly stable
 Most ions represent elements trying to
achieve noble gas electron configurations
 Valence (outer) shell electrons are lost
(cations) or gained (anions) to produce ions.

Cations (atom loses electrons)
Anions (atom gains electrons)
Ion sizes

Anions are normally larger than their
atom “parent”, because they are gaining
electrons

Cations are normally smaller than their
atom “parent” because they are losing
electrons

Anions are generally larger than cations