1. Bohr model and electron configuration

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 Nucleus Electron Electron Orbit Orbit
Energy Levels
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

6. Low energy
High energy
Radiowaves
Microwaves
Infrared .
Ultra-violet
X-Rays
GammaRays
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
Helium
Carbon

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. } Fifth Fourth Increasing energy Third Second First
Further away from the nucleus means more energy.
There is no “in between” energy
Energy Levels
Fifth
Fourth
Third
Increasing energy
Second
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 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
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. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p

24. 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

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

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

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

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

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