Electron
Configuration and
the Periodic Table
Mallard Creek Chemistry - Rines

Electromagnetic
Radiation
Property of Waves
Wave Nature of
Light

 Frequency
▪
No. of waves per
second
 Wave Length
▪
Distance between
corresponding points in
a wave
 Amplitude
▪
Size of the wave peak
Electromagnetic
Radiation

Mathematical Relations
C=λf
C
 This
= speed of light = 3.0 x
108 m/s
 λ (lamda) = wavelength (m)
 f= frequency (Hz or s-1)
is how we know what color light
is emitted!
Frequency is inversely proportional to
Wavelength
λ increases f decreases
If f increases λ decreases
If
Speed of the wave is always constant at 3.0 x 108 m/s
Bohr Model



Nucleus
Energy
Levels

Nucleus: Neutrons and
Protons
Orbits: Electrons
We know both specific
energy and location of
each electron
Electrons orbit the
nucleus in certain fixed
energy levels (or shells)
Bohr Model

Bohr’s Atomic
Model of Hydrogen



Bohr - electrons exist in
energy levels AND
defined orbits around
the nucleus.
Each orbit corresponds
to a different energy
level.
The further out the
orbit, the higher the
energy level
Bohr’s Model

The Photoelectric
Effect
 Light releases electrons
 Not all colors work

Atomic Emission Spectra
 Hydrogen gas emitted
specific bands of light
 Bohr’s calculated energies
matched the IR, visible,
and UV lines for the H
atom
65
4
3
2
1

Electromagnetic
Radiation
Photoelectric Effect – There is a
minimum frequency to eject the
electron
Electromagnetic
Radiation
Only explained by “energy
 Photoelectric

Effect


packets” of light called a
quantum
Quantum - minimum amount of
energy that can be gained or
lost by an atom
Photons are massless particles
of light of a certain quantum of
energy
 Based on the frequency and
wavelength of the photon
Bohr’s Model

Excited electrons
 Energy added to
atom – electrons
“jump” up energy
levels
 When the atom
relaxes - electron
“falls” to lower
energy levels and
emits photon
 Bohr Model of
hydrogen
 Reference Sheets!!!!!

Electromagnetic
Radiation
Atomic Line
Spectra


Electrons in an atom
add energy to go to an
“excited state”.
When they relax back
to the ground state,
they emit energy in
specific energy quanta
Electromagnetic
Radiation

These observations suggested that electrons must
exist in defined energy levels
First, the electron absorbs energy and jumps
from the ground state to an excited state
5
______
4
______
3
______
2
______
1
______
hv
Next, the excited electron relaxes to a
lower excited state or ground state
5
______
5
______
4
______
4
______
3
______
3
______
2
______
2
______
1
______
1
______
hv

Electromagnetic
Radiation
Particle
Wave nature could not
Nature of
Light

explain all observations
(Plank & Einstein)
Photoelectric Effect
E = hf
When light strikes a metal
electrons are ejected
Atomic Line Spectra
▪
When elements are heated,
they emit a unique set of
frequencies of visible and
non-visible light.
Other Scientists
Contributions


De Broglie
Heisenburg




Modeled electrons as waves
Heisenberg Uncertainty
Principle: states one cannot
know the position and
energy of an electron
Electrons exist in orbital’s
of probability
Orbital - the area in space
around the nucleus where
there is a 90% probability
of finding an electron
Other Scientists
Contributions
 Schrödinger


Schrödinger Wave Equation
- mathematical solution of
an electron’s energy in an
atom
Quantum Mechanical Model
of the atom – current
model of the atom treating
electrons as waves.
Quantum Mechanical Model



Nucleus: Neutrons and
protons
Orbitals: region in space
surrounding the nucleus
where there is a 95%
probability of finding an
electron.
We know either energy
or location of each
electron.
Solutions to the Wave
Equation

Quantum
Numbers

Wave Equation
generates 4 variable
solutions
 n - size
 l - shape
 m - orientation
 s – spin

Address of an electron
Quantum Numbers

n – Primary
Quantum Number

Describes the size and
energy of the orbital

n is any positive #

n = 1,2,3,4,….


Found on the periodic
table
Like the “state” you live
in
Quantum Numbers

l – Orbital
Quantum
Number


n=3
l = 0,1,2
n=2
Sub-level of energy
Describes the shape of
the orbital

l = 0,1,2,3,4,….(n-1)

“City” you live in
l = 0,1
n=1
l=0
Quantum Numbers

l – Orbital
Quantum
Number




# level = # sublevels
1st level – 1 sublevel
2nd level – 2 sublevels
4th level = 4 sublevels
Quantum Numbers

s
Sublevels are named for their shape

 l=0
 Spherical in shape

 l=2

p
d
f
 l=3
 l=1
 Dumbbell in shape
s
p
d
f
Quantum Numbers

m – Magnetic
Quantum
Number


Describes the orientation
of the orbital in space
Also denotes how many
orbital's are in each
sublevel

For each sublevel there
are 2l +1 orbital's

“Street” you live on
Quantum Numbers

Look at Orbital's as Quantum Numbers
l=0 m=0
Can only be one
s orbital
l = 1 m = -1, 0, +1
For each p sublevel there are 3
possible orientations, so three 3
orbital's
Orbital Designations
Orbital
Designation
n
l
M
2l+1
No. of No. of
Orbital Electron
3d
3
2
-2,-1,0,+1,+2
5
10
3p
3
1
-1,0,+1
3
6
3s
3
0
0
1
2
2p
2
1
-1,0,+1
3
6
2s
2
0
0
1
2
1s
1
0
0
1
2
Orbital Rules
Number of
Sub-levels
No. of
Orbitals
No. of
Electrons
n
n2
2n2
Energy
Level
Possible
sub-levels
4
s, p, d, f
4
16
32
3
s, p, d
3
9
18
2
s, p
2
4
8
1
s
1
1
2
Reflection


How is the Bohr model different from the
earlier models of the atom?
Who contributed to the modern model of
the atom? How is it different from
Bohr’s?

Why do atoms give unique atomic line
spectra?

What are ground and excited states?

Is 2d possible? 4f ? 2s ? 6p? 1p?
Aufbau Principle

Aufbau
Principal


Lowest energy orbital
available fills first
“Lazy Tenant
Rule”
Pauli’s Exclusion
Principle

Pauli
Exclusion
Principle
 No two electrons have
the same quantum #’s
 Maximum electrons in
any orbital is two ()
Hund’s Rule

Hund’s Rule
RIGHT
 When filling degenerate
orbital's, electrons will fill an
empty orbital before pairing
up with another electron.
 Empty room rule
WRONG
Periodic Table & Electron
Configuration
Periodic Table & Electron
Configuration
Using the periodic table for the filling order of orbitals, by
going in atomic number sequence until you use all the
needed electrons in the element
Increasing Energy
Orbital Energy
Diagram
Sub-level
______ ______ ______ ______ ______
p
______ ______ ______
3 s
Level
(n)
(l)
d
p
______
______ ______ ______
2 s
______
1 s
______
Orbitals (m)
An energy diagram for the
first 3 main energy levels
Orbital Energy Diagram and
Electron Configuration
Increasing Energy
p
3 s
p
______ ______ ______
1s2 2s2 2px2 2py2 2pz2
______
______ ______ ______
2 s
______
1 s
______
Electron
Spin
An energy diagram
for Neon
1s2 2s2 2p6
Electron Configuration
Notation
Orbital Notation


Orbital Notation shows each orbital
O (atomic number 8)
____ ____ ____ ____ ____ ____
1s
2s
2px 2py 2pz 3s
 1s22s22p4
electron configuration!
Orbital Notation


Orbital Notation shows each orbital
O (atomic number 8)
____ ____ ____ ____ ____ ____
1s
!
2s
2px 2py 2pz 3s
Orbital Notation


Write the orbital notation for S
S (atomic number 16)
___ __
1s 2s


__
__
2p
1s22s22p63s23p4
__
__
3s
__ __ __
3p
How many unpaired electrons does sulfur
have?
2 unpaired electrons!
Valence Electrons

Valence Electrons



As (atomic number 33)
 1s22s22p63s23p64s23d104p3

The electrons in
the outermost
energy level.
s and p electrons
in last shell
5 valence
electrons
Valence Electrons

Longhand Configuration
S 16e- 1s2 2s2 2p6 3s2 3p4
Core Electrons
Valence Electrons
Shorthand Configuration
S
16e
[Ne]
2
3s
4
3p

Noble Gas
Configuration
Example - Germanium
1
2
3
4
5
6
7
X X X
X X
X X X
X X
X X
X
[Ar]4s2 3d10 4p2
Electron
Configuration
Let’s Practice
Noble Gas Configuration

P



1s22s22p63s23p3
Ca
(atomic number 20)
As
(atomic number 33)


(atomic number 15)
1s22s22p63s23p64s2
 1s22s22p63s23p64s23d104p3
[Ne] 3s23p3
[Ar] 4s2
[Ar] 4s23d104p3
W (atomic number 74)
 1s22s22p63s23p64s23d104p65s24d105p66s24f145d4
[Xe] 6s24f145d4
Electron
Configuration
Noble Gas
Your Turn
 N (atomic number 7)


1s22s22p3
Na
(atomic number 11)
 1s22s22p63s1

Configuration
[He] 2s22p3
[Ne] 3s1
Sb (atomic number 51)
 1s22s22p63s23p64s23d104p65s24d105p3

Cr
(atomic number 24)
 1s22s22p63s23p64s23d4
[Kr]5s24d105p3
[Ar] 4s23d4
Stability



Full energy level
Full sublevel
Half full sublevel
1
2
3
4
5
6
7
Exceptions


Copper
 Expect: [Ar] 4s2 3d9
 Actual: [Ar] 4s1 3d10
Silver
 Expect: [Kr] 5s2 4d9
 Actual: [Kr] 5s1 4d10

Chromium

Molybdenum
 Expect: [Ar] 4s2 3d4
 Actual: [Ar] 4s1 3d5
 Expect: [Kr] 5s2 4d4
 Actual: [Kr] 5s1 4d5
Exceptions are
explained, but not
predicted!
Atoms are more
stable with half full
sublevel
Stability


Atoms create stability by losing, gaining or
sharing electrons to obtain a full octet
Isoelectronic with noble gases
0
+1
1
2
3
4
5
6
7
+2
+3 +4 -3 -2 -1
Atoms take electron configuration of the closest
noble gas
Stability

Na (atomic number 11)
 1s22s22p63s1
 1s22s22p6 = [Ne]
1
2
3 Na
4
5
6
7
1 Valence electron
Metal = Loses
Ne
Try Some

P-3 (atomic number 15)
 1s22s22p63s23p6

Ca+2 (atomic number 20)


1s22s22p63s23p6
Zn+2 (atomic number 30)
Full Octet
 1s22s22p63s23p63d10
 Lost valence electrons (s and p)
Lewis Structures

Shows valence electrons only!
 s & p electrons
1.
Write noble gas configuration for the element
2.
Place valence electrons around element symbol in order
p electrons
X
4
6
3
7
5
8
1
2
s electrons
Try Some


Write the Lewis structures for:
••
Oxygen (O)
• ••
O
•
– [He] 2s2 2p4

Iron (Fe)
– [Ar] 4s2 3d6

Valence
electrons
Bromine (Br)
– [Ar] 4s2 3d10 4p5
Fe
••
Br
•
••
•
•
••
What Do I Need to
Know?





How the periodic table is arranged
Be able to identify subcategories of the
periodic table
How the elements within a group are
similar
How the elements within a period are
similar
Be able to compare and contrast the
electronegativities, ionization energies,
and radii of metals and non-metals
Periodic Table
Dmitri Mendeleev – Father of the Periodic Table

What He Did
Put elements in rows
by increasing atomic
weight



Put elements in
columns by similar
properties
Some Problems
He left blank spaces for
what he said were
undiscovered elements
(he was right!)
He broke the pattern of
increasing atomic weight
to keep similar reacting
elements together
Mosley
Arranged by Atomic #
Columns = Groups
Rows = Periods
Periodic Table
Organization
Metalloids
Metals
Non-Metals
Periodic Table
Organization
Representative Elements
Transition Metals
Inner Transition Metals
Metals and Nonmetals
Metals








Shiny
Malleable
Ductile (pulled into
wires)
Conduct heat and
electricity
Low specific heat
High melting points
Solids
Lose electrons
Non-metals






Dull
Brittle
Poor conductors
Low melting/boiling
points
Varied properties
Varied phases
Atomic Radius

Atomic Radius = ½ the distance between
adjacent nuclei
 Increases towards Francium
Ionic Radius

Cations
Positive Ion
Metals
Lose electrons

Radius gets smaller!


K

Anions
Negative Ion
Non-metals
Gain electrons

Radius gets larger!


K+
Cl
Cl
Ionization Energy


Energy required to remove an electron from an
atom
Why are there peaks in this trend?
Ionization Energy
Noble gases have the highest first
Ionization Energy
Electronegativity



Pull of electrons in a covalent bond
“Attraction” of atoms towards an electron
Fluorine is “the man”
Periodic Trends
Orbital Size increases
Atomic radius increases
Ionization energy decreases
Electronegativity decreases
Nuclear Charge increases
Atomic radius decreases
Ionization energy increases
Electronegativity increases
What Do I Need To
Know?






How are electrons arranged in an atom
The two natures of electromagnetic
radiation: Particles vs. Waves
How to use the periodic table to list the
configuration or orbital diagram
What quantum numbers are and how they
are related to electron configuration.
How the periodic table is arranged
The basic periodic trends