Mrs. A. Kay
Chem 11

Why don’t the electrons fall into the nucleus?
Move like planets around the sun.
In circular orbits at different levels.
Amounts of energy separate one level from another.

Fixed energy related to the orbit
Electrons cannot exist between orbits
The higher the energy level, the further it is away from the nucleus
An atom with maximum number of electrons in the outermost orbital energy level is stable (unreactive)

He used mathematics to explain the visible spectrum of hydrogen gas http://www.mhhe.com/physsci/chemistr y/essentialchemistry/flash/linesp16.swf

Low energy
Radio waves
Low
Micro waves
Frequency
Long
Wavelength
High energy
.
Infrared Ultraviolet
X-
Rays
Gamma
Rays
High
Frequency
Short
Wavelength
Visible Light

electricity passed through a gaseous element emits light at a certain wavelength
Can be seen when passed through a prism
Every gas has a unique pattern
(color)

Line spectrum of various elements

His theory helped to explain periodic law
Halogens are so reactive because it has one e- less than a full outer orbital
Alkali metals are also reactive because they have only one e- in outer orbital

Bohr’s theory did not explain or show the shape or the path traveled by the electrons.
His theory could only explain hydrogen and not the more complex atoms

Fifth
Fourth
Third
Second
First
Further away from the nucleus means more energy.
There is no
“in between” energy
Energy Levels

The Quantum Mechanical Model
Energy is quantized. It comes in chunks.
A quanta is the amount of energy needed to move from one energy level to another.
Since the energy of an atom is never “in between” there must be a quantum leap in energy.
Schrödinger derived an equation that described the energy and position of the electrons in an atom

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

1 s orbital for every energy level
1s 2s 3s
Spherical shaped
Each s orbital can hold 2 electrons
Called the 1s, 2s, 3s, etc.. orbitals

Start at the second energy level
3 different directions
3 different shapes
Each orbital can hold 2 electrons

The p Sublevel has
3 p orbitals

The D sublevel contains 5 D orbitals
The D sublevel starts in the 3 rd energy level
5 different shapes (orbitals)
Each orbital can hold 2 electrons

The F sublevel has 7 F orbitals
The F sublevel starts in the fourth energy level
The F sublevel has seven different shapes (orbitals)
2 electrons per orbital

Sublevel s p f d
# of shapes
(orbitals)
Max # of electrons
1
3
5
7
2
6
10
14
Starts at energy level
1
2
4
3

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

First Energy Level only s sublevel (1 s orbital) only 2 electrons
1s 2
Second Energy Level s and p sublevels (s and p orbitals are available)
2 in s, 6 in p
2s 2 2p 6
8 total electrons

Third energy level s, p, and d orbitals
2 in s, 6 in p, and 10 in d
3s 2 3p 6 3d 10
18 total electrons
Fourth energy level s,p,d, and f orbitals
2 in s, 6 in p, 10 in d, and 14 in f
4s 2 4p 6 4d 10 4f 14
32 total electrons

7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
3p
2p
6d
5d
4d
3d
5f
4f

Hund’s Rule - When electrons occupy orbitals of equal energy they don’t pair up until they have to .

7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
3p
2p
6d
5d
4d
5f
4f
3d
The first to electrons go into the 1s orbital
Notice the opposite spins only 13 more

7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
3p
2p
6d
5d
4d
3d
The next electrons go into the 2s orbital only 11 more
5f
4f

7s
6s
5s
4s
3s
2s
1s
7p
6p
6d
5f
5d
5p
4d
4p
3d
3p
2p
• The next electrons go into the 2p orbital
• only 5 more
4f

7s
6s
5s
4s
3s
2s
1s
7p
6p
6d
5f
5d
5p
4d
4p
3d
3p
2p
• The next electrons go into the 3s orbital
• only 3 more
4f

7s
6s
5s
4s
3s
2s
1s
7p
6p
6d
5f
5d
5p
4d
4p
3p
2p
3d
• The last three electrons go into the 3p orbitals.
• They each go into separate shapes
• 3 unpaired electrons
• 1s 2
2s
2
2p
6
3s
2
3p
3
4f

Lowest energy to higher energy.
Adding electrons can change the energy of the orbital.
Half filled orbitals have a lower energy.
Makes them more stable.
Changes the filling order

Titanium - 22 electrons
1s
2
2s
2
2p
6
3s
2
3p
6
4s
2
3d
2
Vanadium - 23 electrons
1s
2
2s
2
2p
6
3s
2
3p
6
4s
2
3d
3
Chromium - 24 electrons
1s
2
2s
2
2p
6
3s
2
3p
6
4s
2
3d
4 is expected
But this is wrong!!

1s
2
2s
2
2p
6
3s
2
3p
6
4s
1
3d
5
Why?
This gives us two half filled orbitals.
Slightly lower in energy.
The same principal applies to copper.

Copper has 29 electrons so we expect
1s
2
2s
2
2p
6
3s
2
3p
6
4s
2
3d
9
But the actual configuration is
1s
2
2s
2
2p
6
3s
2
3p
6
4s
1
3d
10
This gives one filled orbital and one half filled orbital.
Remember these exceptions

Time to practice on your own filling up electron configurations.
Do electron configurations for the first
20 elements on the periodic table.