1. Bohr model and electron
configuration
Mrs. A. Kay
Chem 11

2. Bohr’s Model
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.

3. Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels

4. Bohr postulated that:
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)

5. How did he develop his theory?
He used mathematics to explain the
visible spectrum of hydrogen gas
http://www.mhhe.com/physsci/chemistr
y/essentialchemistry/flash/linesp16.swf

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

7. The line spectrum
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)

8. Line spectrum of various elements

9. Bohr’s Triumph
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

10. Drawback
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

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

12. 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

13. 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

14. S orbitals
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

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

16. The p Sublevel has
3 p orbitals

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

18. 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

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

20. 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

21. Electron Configurations
First Energy Level
only s sublevel (1 s orbital)
only 2 electrons
1s2
Second Energy Level
s and p sublevels (s and p orbitals are
available)
2 in s, 6 in p
2s22p6
8 total electrons

22. Third energy level
s, p, and d orbitals
2 in s, 6 in p, and 10 in d
3s23p63d10
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
4s24p64d104f14
32 total electrons

23. 7p 6d
7s 6p 5f
5d
6s 5p 4f
4d
5s
Increasing energy
4p
4s 3d
3p
3s
2p
2s
1s

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

25. 7p 6d
7s 6p 5f
5d
6s 5p 4f
4d
5s
Increasing energy
4p
4s 3d
3p The first to electrons go
3s
into the 1s orbital
2p
2s Notice the opposite
spins
only 13 more
1s

26. 7p 6d
7s 6p 5f
5d
6s 5p 4f
4d
5s
Increasing energy
4p
4s 3d
3p The next electrons go
3s
into the 2s orbital
2p
2s only 11 more
1s

27. 7p 6d
7s 6p 5f
5d
6s 5p 4f
4d
5s
Increasing energy
4p
4s 3d
3p • The next electrons go
3s
into the 2p orbital
2p
2s • only 5 more
1s

28. 7p 6d
7s 6p 5f
5d
6s 5p 4f
4d
5s
Increasing energy
4p
4s 3d
3p • The next electrons go
3s
into the 3s orbital
2p
2s • only 3 more
1s

29. 7p 6d
7s 6p 5f
5d
6s 5p 4f
4d
5s
Increasing energy
4p
4s 3d
3p • The last three electrons
3s go into the 3p orbitals.
2p • They each go into
2s separate shapes
• 3 unpaired electrons
1s • 1s22s22p63s23p3

30. Orbitals fill in order
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

31. Write these electron
configurations
Titanium - 22 electrons
1s22s22p63s23p64s23d2
Vanadium - 23 electrons
1s22s22p63s23p64s23d3
Chromium - 24 electrons
1s22s22p63s23p64s23d4 is expected
But this is wrong!!

32. Chromium is actually
1s22s22p63s23p64s13d5
Why?
This gives us two half filled orbitals.
Slightly lower in energy.
The same principal applies to copper.

33. 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

34. Great site to practice and
instantly see results for
electron configuration.

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