Something Smaller
Than An Atom?
Atomic Structure
Review of the Atom
Smallest part of matter representing an
element
 Once was thought to be the smallest part
of matter
 Later, scientists discovered atoms are
made of subatomic particles

Subatomic particles
Protons - positive charge
 Neutrons – no charge (neutral)
 Electrons – negative charge

Representing a Specific Element
Mass Number
(P + + No)
A
Atomic Symbol
X
Atomic
Number
(P+)
Z
Nucleus: Center Stage

Ernest Rutherford discovered atoms have
a nucleus (1911)
Later scientists discovered that the
nucleus also contains all the neutrons
Rutherford’s Model



Discovered the
nucleus
Electrons moved
around in Electron
cloud
Mostly empty space
Where are those electrons?
Early models of the atom showed
electrons spinning around the nucleus
randomly
 Research showed that this is NOT true

Bohr’s Model
Why don’t the electrons fall into the
nucleus?
 Move like planets around the sun.
 In circular orbits at different levels.
 Energy separates one level from
another.

Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
More about Bohr…..
Created a model showing electrons are in
orbits of different energy around the
nucleus
 …… Think of the planets orbiting the sun

Bohr used the term energy
levels (or shells) to describe
these orbits with different
amounts of energy. He also said that the energy of an
electron is quantized, meaning
electrons can have one energy level
or another but nothing in between.
Bohr’s Model
Increasing energy
Fifth

Fourth
Third
Second

First

Nucleus
Further away
from the nucleus
means more
energy.
There is no “in
between” energy
Energy Levels
Let’s get close to the nucleus
Bohr found that the closer an electron is to
the nucleus, the less energy it needs.
 The farther away it is, the more energy it
needs
 He numbered the electron’s energy levels

Energy levels
1st level can hold up to 2 electrons
 2nd level can hold up to 8 electrons
 3rd level can hold up to 18 electrons
 And so on……..

Energy levels
E-level an electron normally occupies is
called ground state
 But, electrons can move to a higher –
energy, less-stable level, or shell, by
absorbing energy.

This higher energy, less-stable state is
called the electron’s excited state.
The electron gets crunk…..

When the electron is finished being
excited it goes back to its ground state by
releasing some of the energy it has
absorbed
Line spectrum…..what is that???

Energy released by electrons sometimes
occupies part of the electromagnetic
spectrum that humans detect as visible
light
Problem with Bohr’s model

Unexplainable observations on complex atoms
until the quantum theory was created
Quantum theory ----- matter has
properties associated with waves.
It is impossible to know the exact position and
momentum (speed and direction) of an electron at
the same time. ------UNCERTAINTY PRINCIPLE
The Quantum Mechanical
Model




Energy is quantized. It comes in chunks.
Quanta - the amount of energy needed to
move from one energy level to another.
Quantum leap in energy.
Treated electrons as waves



The Quantum Mechanical
Model
Does have energy levels
for electrons.
Orbits are not circular.
It can only tell us the
probability of finding
an electron a certain distance from the
nucleus.


The Quantum Mechanical
Model
The electron is found
inside a blurry
“electron cloud”
A area where there is
a chance of finding an
electron.
Quantum mechanical model of the
atom
Pay attention so you don’t get LOST
!!!!!!!
Atomic Orbitals




Principal Quantum Number (n) = the
energy level of the electron.
Within each energy level the complex
math of Schrödinger's equation describes
several shapes.
These are called atomic orbitals
Regions where there is a high probability
of finding an electron.
S orbitals




1 s orbital for every energy level
Spherical
shaped
Each s orbital can hold 2 electrons
Called the 1s, 2s, 3s, etc.. orbitals.
P orbitals




Start at the second energy level
3 different directions
3 different shapes (dumbell)
Each can hold 2 electrons
D orbitals



Start at the third energy level
5 different shapes
Each can hold 2 electrons
F orbitals
Start at the fourth energy level
 Have seven different shapes
 2 electrons per shape

Summary
# of
Max
shapes electrons
Starts at
energy level
s
1
2
1
p
3
6
2
d
5
10
3
f
7
14
4
Bed check for Electrons…Electron
Configurations……
Goal: Use an energy level diagram to
depict electrons for any element
 1s orbital is closest to the nucleus and it
has the lowest energy
 At energy level 2, there are both s and p
orbitals, with the 2s having lower energy
than the 2p.

Energy level diagram continued
The three 2p subshells are represented by
three dashes of the same energy.
 Energy levels 3, 4, and 5 are also shown
 Notice, that 4s has lower energy than the
3d. This is an exception to what you thought but it

does occur in nature like this.
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
6d
5d
4d
3d
5f
4f
Electron Configurations
The way electrons are arranged in
atoms.
 Aufbau principle- electrons enter the
lowest energy first.
 This causes difficulties because of the
overlap of orbitals of different energies.
 Pauli Exclusion Principle- at most 2
electrons per orbital - different spins

Aufbau Principle

Method for remembering the order in which orbitals fill
the vacant energy levels (like a people fill up vacant
rooms in a hotel)
RULE: ELECTRONS FILL THE LOWEST VACANY
ENERGY LEVELS FIRST.
ANOTHER RULE: WHEN THERE’S MORE THAN ONE
SUB-SHELL AT A PARTICULAR ENERGY LEVEL, SUCH AS
AT THE 3P OR 4 D LEVELS, ONLY ONE ELECTRON FILLS
EACH SUB-SHELL UNTIL EACH SUBSHELL HAS ONE
ELECTRON.
THEN, ELECTRONS START PAIRING UP IN EACH
SUBSHELL.
THIS RULE IS CALLED HUND’S RULE.
Electron Configuration
Hund’s Rule- When electrons occupy
orbitals of equal energy they don’t pair up
until they have to .
 Let’s determine the electron configuration
for Phosphorus
 Need to account for 15 electrons

Increasing energy
7s
6s
5s
7p
6p
6d
5d
5p
4d
4p
3s
2p
2s



1s
4f
3d
4s
3p
5f
The first to electrons go
into the 1s orbital
Notice the opposite
spins
only 13 more
Increasing energy
7s
6s
5s
7p
6p
6d
5d
5p
4d
4p
3s
2p
2s
1s
4f
3d
4s
3p
5f


The next electrons go
into the 2s orbital
only 11 more
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p • The next electrons
go into the 2p orbital
2p
• only 5 more
5f
4f
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p • The next electrons
go into the 3s orbital
2p
• only 3 more
5f
4f
Increasing energy
7s
6s
5s
4s
7p
6p
6d
5d
5p
4d
4p
3p •
3s
2s
1s
2p •
•
•
5f
4f
3d
The last three electrons
go into the 3p orbitals.
They each go into
separate shapes
3 unpaired electrons
1s22s22p63s23p3
The easy way to remember
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
1s
• 2 electrons
Fill from the bottom up
following the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
1s 2s
• 4 electrons
Fill from the bottom up
following the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
• 12 electrons
Fill from the bottom up
following the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
6
2
3p 4s
• 20 electrons
Draw the energy level diagram of
oxygen
Find oxygen on the period table
 Atomic number = 8
 Therefore, it has 8 protons and 8 electrons
 So, you put 8 electrons into your energy
level diagrams

Energy level diagram for oxygen
Use arrows to represent electrons
 If two electrons fit in the same orbital, one
must face up and the other one down

The first electron goes into the 1s orbital, filling the lowest energy level
first.
And, the second one spin pairs with the first one.
Electrons 3 and 4 spin pair in the next lowest vacant orbital – the 2 s.
Electron 5 goes into one of the 2p sub-shells
Electrons 6 and 7 go into the other two totally vacant 2p orbitals
The last electron spin pairs with one of the electrons in the 2p subshells.
Electron Configuration

Oxygen 1s22s22p4

You can derive the electron configuration from the energy level
diagram.
The first two electrons in oxygen fill the 1s orbital, so you it as 1s2 in
the electron configuration.


The 1 is the energy level, the s represents the type of orbital, and
the superscript 2 represents the number of electrons in that orbital.

The next two electrons are in the 2s orbital, so you write 2s2.

Last, you show the 4 electrons in the 2p orbital as 2p4.

That’s how you get 1s22s22p4.
Tip

The sum of the superscript numbers
equals the atomic number, or the number
of electrons in the atom.
Exceptions to
Electron
Configuration
Orbitals fill in order
Lowest energy to higher energy.
 Adding electrons can change the energy
of the orbital.
 Filled and half-filled orbitals have a lower
energy.
 Makes them more stable.
 Changes the filling order of d orbitals

Copper’s electron configuration
Copper has 29 electrons so we expect
 1s22s22p63s23p64s23d9
 But the actual configuration is
 1s22s22p63s23p64s13d10
 This gives one filled orbital and one half
filled orbital.
 Remember these exceptions

 s 2d 4  s 1d 5
 s2d9  s1d10
Here are a couple of electron configurations you can use to
check your conversions from energy level diagrams:
Chlorine (Cl)
Iron (Fe)
Answers
Answer: Cl
1s22s22p63s23p5
Answer: Fe 1s22s22p63s23p64s23d6