Bohr model and electron configuration

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
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. Radio
waves
Micro
waves
Infrared
.
Ultra-
violet
X-
Rays
Gamma
Rays
Low
energy
High
energy
Low
Frequency
High
Frequency
Long
Wavelength
Short
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
from the
nucleus
means more
energy.
There is no
“in between”
energy
Energy Levels
First
Second
Third
Fourth
Fifth
Increasing energy

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 2s3s
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 3
rd
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. Summarys
p
d
f
# of shapes
(orbitals)
Max # of
electrons
1 2 1
3 6 2
5 10 3
7 14 4
Sublevel
Starts at
energy
level

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
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

22. 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

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

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

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

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

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

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

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

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
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!!

32. Chromium is actually
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.

33. Copper’s electron configuration
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

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.